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The document discusses elementary substances, the periodic table of elements, and the characteristics of atoms, including protons, neutrons, and electrons. It classifies elements into metals, nonmetals, and metalloids, detailing their properties, behaviors, and group classifications based on electron configurations. Additionally, the text explains isotopes and provides methods for calculating atomic mass.

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ELEMENTARY SUBSTANCES ELEMENTS AND THE PERIODIC TABLE OF ELEMENTS ATOMS: THE SMALLEST PARTICLE OF MATTER EVIDENCES OF SUBATOMIC PARTICLES DETERMINATION OF ATOMIC MASSES THE ELECTRONIC CONFIGURATION BY MAIN ENERGY LEVEL AND BY ENERGY SUBLEVEL PERIODICITY AND PERIODIC RELATIONSHIPS
ELEMENT A PURE SUBSTANCE WHICH IS COMPOSED OF A SINGLE TYPE OF ATOM IT IS CONSISTS ONLY OF ATOMS THAT ALL HAVE THE SAME NUMBER OF PROTONS IN THEIR ATOMIC NUCLEI THE NUMBER OF PROTONS IN THE NUCLEUS IS THE DEFINING PROPERTY OF AN ELEMENT, AND IS REFERRED TO AS ITS ATOMIC NUMBER (z). ALL ATOMS WITH THE SAME ATOMIC NUMBER ARE ATOMS OF THE SAME ELEMENT IT CANNOT BE BROKEN DOWN INTO SIMPLER SUBSTANCES BY CHEMICAL MEANS. WHEN DIFFERENT ELEMENTS UNDERGO CHEMICAL REACTION, ATOMS ARE REARRANGED INTO NEW COMPOUNDS HELD TOGETHER BY CHEMICAL BONDS. 13p 14n 13e
ELEMENTS CAN BE EXPRESSED USING CHEMICAL SYMBOLS. BORON GOLD TITANIUM SULFUR TANTALUM BARIUM URANIUM TITANIUM SULFUR TANTALUM
EXERCISES. DECIPHER THE FOLLOWING USING CHEMICAL SYMBOLS SODIUM NEON THORIUM EINSTEINIUM PROTACTINI UM RHENIUM SULFUR PRASEODYM IUM INDIUM CERIUM SULFUR LANTHANUM URANIUM RHENIUM NITROGEN TELLURIUM MOLYBDENU M NICKEL COBALT MOLYBDENU M NITROGEN IODINE CARBON OXYGEN CARBON AMERICIUM POLONIUM IRIDIUM IODINE SULFUR NICKEL IODINE POTASSIUM POTASSIUM IODINE SULFUR AMERICIUM OXYGEN NITROGEN TELLURIUM PROTACTINI UM RADIUM IODINE SULFUR OXYGEN SODIUM, CARBON, IODINE, OXYGEN,
TYPES OF ELEMENTARY SUBSTANCES
B B B B B B B B B B B B B B B B B B B B B B MET ALS MET ALS NON MET ALS B
METALLIC ELEMENTS Nearly all are shiny and grey-white in color Cu, Cs, and Au are shiny and golden Reflectivity is from intermediate to typically high. Almost all are solid in form. Rb, Cs, Fr, Cs and Hg are liquid at/near STP Densities are generally high. with some exceptions such as alkali metals Most metals are ductile and malleable. Some are brittle like Cr, Mn, Ge, Ru, W, Os, and Bi Most have hexagonal and cubic structure at freezing point. Close-packed crystal structures with high coordination numbers Atomic radius is from intermediate to very large. 112–298 pm (average 187 pm) Around half form allotropes Sn has a metalloid-like allotrope (grey Sn) which forms below 13.2 °C. Thermal conductivity is from medium to high Melting point is mostly high and generally expands during melting. Enthalpy of fusion is from low to high.
s, p, d, and f periodic table block. Few number of electrons in outer s and p. except for Pd (0), Sn, Pb, and Fl (with 4), Bi (with 5) and Po (with 6) All has free electrons that facilitates electrical and thermal conductivity. Electrical conductivity is from good to high but falls gradually as temperature rises (as it turns to liquid) Over-all metallic behavior All metals tend to form cations. Metals seldom form covalent compounds (non-covalent bond) Oxidation number is nearly always positive Ionization energy is relatively low. Electronegativity is usually low. They are base-forming Their names usually end in –um or –ium with the following exceptions: Bi, W, Mn, Ni, and Co.
nonMETALLI C ELEMENTS most are colourless or dull red, yellow, green, or intermediate shades C, P, Se, and I are shiny and grey-white Reflectivity is from zero or low (mostly) to intermediate. most are gases. C, P, S, Se, and I are solid while Br is liquid. Their densities are often low. They are brittle when solid Some (C, P, S, and Se) have non-brittle forms. Their crystalline structure at freezing point ranges from hexagonal (H, He, C, N, and Xe), to cubic (O, F, Ne, P, Ar, Kr, Xe, and Rn) and orthorhombic ( S, Cl, Br & I). They have open structures & low coordination numbers. Atomic radius is from very small to intermediate 31–120 pm, average 76.4 pm
s and p periodic table block. Most high number of electrons in the outer s and p except H and He. no, few, or directionally confined free electrons (generally hampering electrical and thermal conductivity) Electrical conductivity is from poor to good but increases as temperature rises Over-all nonmetallic behavior All nonmetals tend to form anions. Nonmetals form many covalent compounds Oxidation number is either positive or negative. Ionization energy and electronegativity are usually high. Some form allotropes while some (like graphitic C, black P and grey Se) are more metalloidal or metallic in nature Thermal conductivity is from almost negligible to very high. They are acid-forming Their names usually end in –on, -gen, or –ine with the following exceptions: S ends with –ur and Te and He end with -ium
METALLOID S shiny and grey- white in color Reflectivity is intermediate. All are solid in form Their densities are lower than nearby metals but higher than nearby nonmetals Their crystalline structure at freezing point ranges from rhombohedral (B, As, and Sb), to cubic (Si and Ge) and hexagonal (Te). relatively open crystal structures and medium coordination numbers[ Atomic radius is from small to intermediate: B, Si Ge, As, Sb, and Te 87–123 pm, average 115.5 pm They are brittle. all or nearly all form allotropes Some (red B and yellow As) are more nonmetallic in nature Si has high thermal conductivity but mostly are intermediate. Melting point is high and some contract, unlike (most) metals.
p periodic table block. medium number (3–7) in the outer level valence electrons less freely delocalized; considerable covalent bonding present Electrical conductivity is from intermediate to good. Most behaves like metals. Over-all nonmetallic behavior Some have tendency to form anions in water. Solution chemistry dominated by formation and reactions of oxyanions. Form salts as well as covalent compounds Oxidation number is either positive or negative. Intermediate ionization energy and electronegativity is close to 2 (between 1.9 and 2.2)
GIVE THE CHEMICAL NAMES AND IDENTIFY WHETHER METAL, NONMETAL OR METALLOID 1. C 2. Ca 3. Cl 4. Ce 5. Cf 6. Cm 7. Co 8. Cr 9. Cs 10. Cu 11. S 12. Sb 13. Sc 14. Sm 15. Sn 16. Sr 17. Mg 18. Mn 19. Mo 20. Mt
DIFFERENT GROUPS OF ELEMENTARY SUBSTANCES
An element family is a set of elements sharing common properties. Elements are classified into families because the three main categories of elements (metals, nonmetals and metalloids) are very broad. The characteristics of the elements in these families are determined primarily by the number of electrons in the outer energy shell. Element groups, on the other hand, are collections of elements categorized according to similar properties. Because element properties are largely determined by the behavior of valence electrons,
PERIODS & GROUPS
PERIODS The rows of the table are known as periods. It is in the succesive periods that we observe the periodicity of properties of the elements. Each period has the full range of properties. For example more metallic elements occur to the left of a period, and the less metallic elements to the right; or oxides of the elements to the left are basic and acidic for elements to the right. The periods are simply numbered 1 though 7 from the top down GROUPS The columns of the table are known as groups or families. All the elements in a group have similar properties. Placing elements in groups is one of the most important ways of classifying them. There is some variation in properties within a group, but the changes are relatively small as one
PERIOD 1 TO 7 GROUP 1 TO 18 IA, IIA, IIIB, IVB, VB, VIB, VIIB, VIIIB, IIIA- VIIIA
B B B B B B B B B B B B B B B B B B FAMI LY A FAMIL Y B
The periodic table groups are as follows (in the brackets are shown the old systems: European and American): Group 1 (IA): the alkali metals Group 2 (IIA): the alkaline earth metals Group 3 (IIIB) Group 4 (IVB) Group 5 (VB) Group 6 (VIB) Group 7 (VIIB) Group 8 (VIIIB) Group 9 (VIIIB) Group 10 (VIIIB) Group 11 (IB): the coinage metals (not a IUPAC-recommended name) Group 12 (IIB) Group 13 (IIIA): the boron group Group 14 (IVA): the carbon group Group 15 (VA): the pnictogens (not a IUPAC-recommended name) or nitrogen group Group 16 (VIA): the chalcogens
ALKALI METALS are recognized as a group and family of elements. These elements are metals. Hydrogen is not considered an alkali metal because the gas does not exhibit the typical properties of the group. However, under the right conditions of temperature and pressure, hydrogen can be an alkali metal. 1 valence electron Soft metallic solids Shiny, lustrous High thermal and electrical conductivity Low densities, increasing with atomic mass Relatively low melting points, decreasing with atomic mass Vigorous exothermic reaction with water to GROU P IA (CAS) GROU P 1 (IUPAC)
ALKALINE EARTH METALS or simply alkaline earths are recognized as an important group and family of elements. These elements are metals. Examples include calcium and magnesium. 2 valence electrons Metallic solids, harder than the alkali metals Shiny, lustrous, oxidize easily High thermal and electrical conductivity More dense than the alkali metals Higher melting points than alkali metals Exothermic reaction with water, increasing as you move down the group; beryllium does not react with water; magnesium reacts only with steam GROUP IIA (CAS) GROUP 2 (IUPAC)
TRANSITION METALS is the largest family of elements. The center of the periodic table contains the transition metals, plus the two rows below the body of the table (lanthanides and actinides) are special transition metals. Groups 3-12 The d and f block metals have 2 valence electrons Hard metallic solids Shiny, lustrous High thermal and electrical conductivity Dense High melting points Large atoms exhibit GROUP IIIB GROUP 3 GROUP VIIIB GROUP 8, 9, 10 GROUP IB / IIB GROUP 11, 12
BORON GROUP or earth metal family is not as well-known as some of the other element families. 3 valence electrons Diverse properties, intermediate between those of metals and nonmetals CARBON GROUP is made up of elements called tetrels, which refers to their ability to carry a charge of 4. 4 valence electrons Diverse properties, intermediate between those of metals and nonmetals GROUP IIIA GROUP 13 GROUP IVA GROUP 14
NITROGEN GROUP or pnictogens is a significant element family. 5 valence electrons Diverse properties, intermediate between those of metals and nonmetals OXYGEN GROUP or the chalcogens family. 6 valence electrons Diverse properties, changing from nonmetallic to metallic in moving down the family GROUP VA (CAS) GROUP 15 (IUPAC) GROUP VIA (CAS) GROUP
HALOGEN FAMILY is a group of reactive nonmetals. 7 valence electrons Reactive nonmetals Melting points and boiling points increase with increasing atomic number High electron affinities Change state as it moves down the family, with fluorine and chlorine existing as gases at room temperature while bromine is a liquid and iodine is a solid NOBLE GASES are a family of nonreactive nonmetals. 8 valence electrons Typically exist as monatomic gases, although these elements do (rarely) form compounds Stable electron octet makes nonreactive (inert) under GROUP VIIA (CAS) GROUP 17 (IUPAC) GROUP VIIIA (CAS) GROUP 18 (IUPAC)
ATOMS • incredibly tiny • numerous • perpetually in motion • ageless • An atom takes part in chemical reactions independently. • An atom can be divided into a number of sub-atomic particles - electron, proton and neutron.
SUB ATOMIC PARTICLES
• ELECTRONS – discovered by Sir John Joseph Thomson in 1897 – are located in an electron cloud, which is the area surrounding the nucleus of the atom – can abbreviated as e- – have a negative charge that is equal in magnitude to the positive charge of the protons – mass is considerably less than that of a proton or neutron
• PROTONS – discovered by Ernest Rutherford in the year 1919, when he performed his gold foil experiment. – exist in a nucleus and have a positive nuclear charge. – atomic number or proton number is the number of protons present in an atom – The atomic number determines an element (e.g., the element of atomic number 6 is carbon).
• NEUTRONS – were discovered by James Chadwick in 1932, when he demonstrated that penetrating radiation incorporated beams of neutral particles – located in the nucleus with the protons – Along with protons, they make up almost all of the mass of the atom – The number of neutrons is called the neutron number and can be found by subtracting the proton number from the atomic mass number.
• u is the SI symbol for atomic mass unit. • The positive charge of protons cancels the negative charge of the electrons. Neutrons have no charge. • With regard to mass, protons and neutrons are very similar, and have a much greater mass than electrons. Compared with neutrons and protons, the mass of an electron is usually negligible. • Spin is associated with the rotation of a particle. Protons, neutrons, and electrons each have a total spin of 1/2.
RELATIONSHIP OF THE ATOMIC MASS AND ATOMIC NUMBER WITH THE SUBATOMIC PARTICLES • A = p + n - Atomic mass (A) is the sum of the number of protons (p) and neutrons (n) in the nucleus of an atom. • Z = p - Atomic number (Z) gives the number of protons (p) in the nucleus. It is what distinguishes the atoms of a given element. • n = A – p - The number of neutrons can be found by subtracting the proton number from the atomic mass number.
Element Atomic Mass (A) Atomic Number (Z) Protons (p) Electrons ( e- ) Neutrons (n) K 39 19 19 19 20 Ar 40 18 18 18 22 Na 23 11 11 11 12 O 16 8 8 8 8 Au 197 79 Ag 108 47 Zn Mg Og Mn I 79 79 118 47 47 61
ISOTOPES: Not all atoms are alike • Isotopes are atoms of the same element with different number of neutrons or different atomic masses • Isotopes are atoms with the same number of neutrons. Isotopes are different forms of a single element.
Known Isotopes of an Element • Hydrogen - 1H - 2H - 3H - 4H - 5H -6H -7H  Lithium - 3Li - 8Li - 4Li - 9Li - 5Li - 10Li - 6Li - 11Li - 7Li - 12Li  Beryllium - 5Be - 12Be - 6Be - 13Be - 7Be - 14Be - 8Be - 15Be - 9Be - 16Be - 10Be - 11Be  Boron - 6B - 12B - 7B - 13B - 8B - 14B - 9B - 15B - 10B - 16B - 11B - 17B  Carbon - 8C - 14C - 9C - 15C - 10C - 16C - 11C - 17C - 12C - 18C - 13C - 19C particle p e- n 1H 1 1 0 2H 1 1 1 3H 1 1 2 particle p e- n 3Li 3 3 0 4Li 3 3 1 5Li 3 3 2 particle p e- n 5Be 4 4 1 6Be 4 4 2 7Be 4 4 3
HOW IS ATOMIC MASS DETERMINED? A = (IM X FA)1 + (IM X FA)2 + (IM X FA)3 + ....... IM is the isotopic mass of a naturally occuring element FA is the isotopic abundance of a naturally occuring element What is the average atomic mass of magnesium? ISOTOPE ISOTOPIC ABUNDANCE ISOTOPIC MASS Mg - 24 78.99 % 23.98501417 u Mg - 25 10.00 % 24.98583692 u Mg - 26 11.01 % 25.98259292 u A = (IM X FA)Mg-24 + (IM X FA)Mg-25 + (IM X FA)Mg-26 A = (23.98501417 X 0.7899) + (24.98583692 X 0.1000) + (25.98259292 X 0.1101) A = 24.30502986 A = 24.31 u
Calculate the atomic mass of potassium given the following: K-39 isotopic abundance- 93.12% isotopic mass- 38.964 u K-41 isotopic abundance- 6.88% isotopic mass- 40.962 u A = (IM X FA)K-39 + (IM X FA)K-41 A = (38.964 X 0. 9312) + (40.962 X 0.0688) A = 39.102 u A = 39.1 u Use the atomic masses of each of the two isotopes of chlorine along with their isotopic abundances to calculate the average atomic mass of chlorine. chlorine-35: atomic mass = 34.969 amu and % abundance = 75.77% chlorine-37: atomic mass = 36.966 amu and % abundance = 24.23% A = (IM X FA)Cl-35 + (IM X FA)Cl-37 A = (34.969 X 0.7577) + (36.966 X 0.2423) A = 35.457 u A = 35.46 u
ELEMENTAL IONS: charged particles • ELEMENTAL IONS ARE PARTICLES THAT RESULTED FROM THE LOSS OR GAIN OF ELECTRON. • WHEN AN ATOM LOSES ELECTRON IT WILL BECOME POSITIVELY CHARGED (CATION). • Na has 11p, 11e-, and 12n. When one of its electron is transferred to another atom, it will become positively charged. The number of p and n are the same but the number of electron becomes 10. The charge of the this ion can be determined by subtracting the number of electrons from the number of proton. Charge = p – e- = 11 – 10 = 1 • Na atom  Na+ ion
• WHEN AN ATOM GAINS ELECTRON IT WILL BECOME NEGATIVELY CHARGED (ANION). • Cl has 17p, 17e-, and 18n. When one electron is accepted by Cl , it will become negatively charged. The number of p and n are the same but the number of electron becomes 18. The charge of the this ion can be determined by subtracting the number of electrons from the number of proton. • Charge = p – e- = 17 – 18 = -1 • Cl atom  Cl- ion
PARTICLE p e- n F fluorine atom 9 9 10 O oxygen atom 8 8 8 N nitrogen atom 7 7 7 S sulfur atom 16 16 16 PARTICLE p e- n F- fluoride ion 9 10 10 O2- oxide ion 8 10 8 N3- nitride ion 7 10 7 S2- sulfide ion 16 18 16
ELECTRON CONFIGUR ATION The electron configuration of an atomic species (neutral or ionic) allows us to understand the shape and energy of its electrons. Many general rules are taken into consideration when assigning the location of the electron to its prospective energy state, however these assignments are arbitrary and it is always uncertain as to which electron is being described. Knowing the electron
ELECTRON CONFIGUR ATION The arrangement or distribution of electrons in shells or
The protons and neutrons are found inside the nucleus of an atom, while the electrons are constantly moving around the nucleus. The nucleus and the electrons interact to form the most stable arrangement possible, the ground state, which requires the least amount of energy. There are four main concepts that must remembered in electron configuration. These are the energy levels, the enrgy sublevel, the orbitals, and number of electrons per sublevel.
Energy Levels and the Atomic Model The concept of energy levels is one part of the atomic model that is based on a mathematical analysis of atomic spectra. Each electron in an atom has an energy signature that is determined by its relationship with other negatively charged electrons in the atom and the positively charged atomic nucleus. An electron can change energy levels, but only by steps or quanta, not continuous increments. The energy of an energy level increases the further out from the nucleus it is. The lower the number of a principal energy level, the closer together the electrons are to each other and to the nucleus of the atom. During chemical reactions, it's more difficult to remove an electron from a lower energy level
Rules of Principal Energy Levels A principal energy level may contain up to 2n2 electrons, with n being the number of each level. The first energy level can contain 2(1)2 or two electrons; the second can contain up to 2(2)2 or eight electrons; the third can contain up to 2(3)2 or 18 electrons, and so on. PRINCIPAL ENERGY LEVEL. n SHELL LETTER MAXIMUM NUMBER OF ELECTRONS, 2n2 1 K 2n2 = 2(1)2 = 2(1) = 2 2 L 2n2 = 2(2)2 = 2(4) = 8 3 M 2n2 = 2(3)2 = 2(9) = 18 4 N 2n2 = 2(4)2 = 2(16) = 32 5 O 2n2 = 2(5)2 = 2(25) = 50 6 P 2n2 = 2(6)2 = 2(36) = 72 7 Q 2n2 = 2(7)2 = 2(49) = 98
ENERGY SUBLEVELS/ENERGY SUBSHELLS Each energy level, or shell, is divided into sublevels. The terms sublevel and subshell are used interchangeably. The sublevels are represented by the letters s, p, d, and f. Each energy level has certain sublevels. ENERGY LEVELS ENERGY SUBLEVELS n = 1 s n = 2 s, p n = 3 s, p, d n = 4 s, p, d, f Energy levels that are higher than four would contain additional sublevels such as g and h. No known atom, in its ground state, would have electrons that
ORBITALS Each sublevel is made up of orbitals. Each sublevel has a different number of orbitals. ENERGY SUBLEVELS ORBITALS s 1 p 3 d 5 f 7 Orbitals are classified by their shape. s orbitals are spherical in shape.
Orbitals found in the p subshell are dumbbell, or figure-eight, shaped. Remember, in each p subshell there are three orbitals, so there are three dumbbell shapes. Atoms are three dimensional objects, and the orbitals are oriented around the different axes. The image below shows the three p orbitals and the complete p sublevel containing all three orbitals.
The first principal energy level has one sublevel that contains one orbital, called the s orbital. The s orbital can contain a maximum of two electrons. The next principal energy level contains one s orbital and three p orbitals. The set of three p orbitals can hold up to six electrons. Thus, the second principal energy level can hold up to eight electrons, two in the s orbital and six in the p orbital. The third principal energy level has one s orbital, three p orbitals, and five d orbitals, which can each hold up to 10 electrons. This allows for a maximum of 18 electrons. The fourth and higher levels have an f sublevel in addition to the s, p, and d orbitals. The f sublevel contains seven f orbitals, which can each hold up to 14 electrons. The total number of electrons in
PRINCIPAL ENERGY LEVEL, n NUMBER OF ORBITALS ORBITALS NUMBER OF ELECTRONS PER ORBITAL 1 1 s 2 2 2 s p 2 6 3 3 s p d 2 6 10 4 4 s p d f 2 6 10 14 Electron Notation The notation used to indicate the type of energy level and the number of electrons in that level has a coefficient for the number of the principal energy level, a letter for the sublevel, and a superscript for the number of electrons located in that sublevel. For example, the notation 4p3 indicates the fourth principal energy level, the p sublevel, and the presence of
An orbital diagram helps to determine the electron configuration of an element. An element’s electron configuration is the arrangement of the electrons in the shells. There are a few guidelines for working out this configuration: 1. Each orbital can hold only two electrons. Electrons that occur together in an orbital are called an electron pair. 2. An electron will always try to enter the orbital with the lowest energy. 3. An electron can occupy an orbital on its own, but it would rather occupy a lower-energy orbital with another electron before occupying a higher- energy orbital. In other words, within one energy
THE BUILDING-UP PRINCIPLE The Aufbau principle determines an atom’s electron configuration by adding electrons to atomic orbitals following a defined set of rules. Although the nucleus of an atom is very dense, the electrons around it can take on a variety of positions which can be summarized as an electron configuration. An element’s electron configuration can be represented using energy level diagrams, or Aufbau diagrams. The Aufbau principle (from the German Aufbau meaning “building up, construction”) describes a model-building method in which an atom is “built up” by progressively
Filling in an Aufbau Diagram The order in which orbitals are filled is given by the Madelung rule. The rule is based on the total number of nodes in the atomic orbital, n + ℓ, which is related to the energy. In this context, n represents the principal quantum number and ℓ represents the azimuthal quantum number. The values ℓ = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. According to the principle, electrons fill orbitals starting at the lowest available energy states before filling higher states (e.g., 1s before 2s). The Madelung rule defines the order in which atomic orbitals are filled with electrons. Electrons fill orbitals starting at the lowest available energy state before filling higher states.
Filling in an Aufbau Diagram
The following steps detail how to draw an Aufbau diagram 1. Determine the number of electrons that the atom has. 2. Fill the s orbital in the first energy level (the 1s orbital) with the first two electrons. 2. Fill the s orbital in the second energy level (the 2s orbital) with the second two electrons. 3. Put one electron in each of the three p orbitals in the second energy level (the 2p orbitals) and then if there are still electrons remaining, go back and place a second
Limitations to Aufbau The Aufbau principle is based on the idea that the order of orbital energies is fixed— both for a given element and between different elements. This assumption is approximately true—enough for the principle to be useful—but not physically reasonable. It models atomic orbitals as “boxes” of fixed energy into which at most two electrons can be placed. However, the energy of an electron in an atomic orbital depends on the energies of all the other electrons of the atom.
Electron Configuration Standard Notation A special type of notation is used to write an atom’s electron configuration. The notation describes the energy levels, orbitals, and the number of electrons in each. For example, the electron configuration of lithium is 1s22s1. The number and letter describe the energy level and orbital, and the number above the orbital shows how many electrons are in that orbital. Using standard notation, the electron configuration of fluorine is 1s22s22p5.
Electrons will fill the lowest energy orbitals first and then move up to higher energy orbitals only after the lower energy orbitals are full.. Although the implications are clear for orbitals of different principal quantum number (n), which are clearly of different energy, the filling order is less clear for degenerate sublevels. For example, for boron through neon, the electron filling order of the 2p orbitals follows Hund’s Rule. Hund’s Rule states that: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied. All of the electrons in singly occupied orbitals have the same spin.
Diamagnetism Any time two electrons share the same orbital, their spin quantum numbers have to be different. In other words, one of the electrons has to be “spin- up,”, while the other electron is “spin-down,” This is important when it comes to determining the total spin in an electron orbital. In order to decide whether electron spins cancel, add their spin quantum numbers together. Whenever two electrons are paired together in an orbital, or their total spin is 0, they are called diamagnetic electrons.
Paramagnetism Electrons that are alone in an orbital are called paramagnetic electrons. Remember that if an electron is alone in an orbital, the orbital has a net spin, because the spin of the lone electron does not get canceled out. If even one orbital has a net spin, the entire atom will have a net spin. Therefore, an atom is considered to be paramagnetic when it contains at least one paramagnetic electron. In other words, an atom could have 10 paired (diamagnetic) electrons,
- Periodic trends are specific patterns that are present in the periodic table that illustrate different aspects of a certain element, including its size and its electronic properties. - Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties. - These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and Periodic Trends
property describing an atom's ability to attract and bind with electrons. Because electronegativity is a qualitative property, there is no standardized method for calculating electronegativity. However, the most common scale for quantifying electronegativity is the Pauling scale, named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature of electronegativity. Electronegativity values for each
From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. If the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one. From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic
Important exceptions of the above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values. As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.
Ionization Energy - is the energy required to remove an electron from a neutral atom in its gaseous phase. Conceptually, ionization energy is the opposite of electronegativity. The lower this energy is, the more readily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely it is that the atom becomes a cation. Generally, elements on the right side of the periodic table have a higher ionization energy because their valence shell is nearly filled. Elements on the left side of the periodic table have low ionization
The ionization energy of the elements within a period generally increases from left to right. This is due to valence shell stability. The ionization energy of the elements within a group generally decreases from top to bottom. This is due to electron shielding. The noble gases possess very high ionization energies because of their full valence shells as indicated in the graph. Note that helium has the highest ionization energy of all the elements.
Electron Affinity - is the ability of an atom to accept an electron. Unlike electronegativity, electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom's affinity for electrons. - Electron affinity increases from left to right within a period. This is caused by the decrease in atomic radius. - Electron affinity decreases from top to bottom within a group. This is caused by the increase in
ATOMIC RADIUS - While the atomic radius can be defined in a number of different ways, the general atomic radius trend across the periodic table holds true. The atomic radius for atoms of an element tends to go up as you move down a group of elements in the table. The atomic radius increases as you move down a column because for every new row of the table a new electron shell is added to the atom. - Atomic radius decreases from left to right within a period. This is caused by the increase in the number of protons and electrons across a period. One proton has a greater effect than one electron; thus, electrons are pulled towards the nucleus, resulting in a smaller radius. Atomic radius increases from top to bottom within a
Melting Point - is the amount of energy required to break a bond(s) to change the solid phase of a substance to a liquid. Generally, the stronger the bond between the atoms of an element, the more energy required to break that bond. Because temperature is directly proportional to energy, a high bond dissociation energy correlates to a high temperature. Melting points are varied and do not generally form a distinguishable trend across the periodic table. Metals generally possess a high melting point. Most non-metals possess low melting points. The non-metal carbon possesses the highest melting point of all the elements. The semi-metal boron also possesses a high melting point.
The Metallic Character - of an element can be defined as how readily an atom can lose an electron. From right to left across a period, metallic character increases because the attraction between valence electrons and the nucleus is weaker, enabling an easier loss of electrons. Metallic character increases as you move down a group because the atomic size is increasing. When the atomic size increases, the outer shells are farther away. The principal quantum number increases and average electron density moves farther
from left to right across a period. This is caused by the decrease in radius of the atom that allows the outer electrons to ionize more readily. Metallic characteristics increase down a group. Electron shielding causes the atomic radius to increase thus the outer electrons ionizes more readily than electrons in smaller atoms. Metallic character relates to the ability to lose electrons, and nonmetallic character relates to the
ELECTRON CONFIGURATION BY MAIN ENERGY LEVEL
ELECTRON CONFIGURATION BY MAIN ENERGY LEVEL THE DISTRIBUTION OR ARRANGEMENT OF ELECTRONS IN THE ATOM’S SHELL. ELECTRONS ARE DISTRIBUTED STARTING FROM THE INNERMOST SHELL. THERE ARE SEVEN SHELLS OR MAIN ENERGY LEVELS WHICH ARE REPRESENTED BY EITHER SHELL LETTER OR SHELL NUMBER.
I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 38. 1
1 1s22s22p63s23p64s23d104p65s2 K SHELL 1s2 2 e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p6 3d10 18 e- N SHELL 4s2 4p6 8 e- O SHELL 5s2 2 e- PERIOD 5 GROUP IIA METAL Sr
Psych e- Fills MEL I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 33. 2
2 1s22s22p63s23p64s23d104p3 K SHELL 1s2 2 e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p6 3d10 18 e- N SHELL 4s2 4p3 5 e- PERIOD 4 GROUP VA NONMETAL As
I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 20. 3
3 1s22s22p63s23p64s2 K SHELL 1s2 2 e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p6 8 e- N SHELL 4s2 2 e- PERIOD 4 GROUP IIA METAL Ca
I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 17. 4
4 1s22s22p63s23p5 K SHELL 1s2 2 e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p5 7 e- PERIOD 3 GROUP VIIA NONMETAL Cl
I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 11. 5
5 1s22s22p63s1 K SHELL 1s2 2 e- L SHELL 2s2 2p6 8 e- M SHELL 3s1 1 e- PERIOD 3 GROUP IA METAL Na
I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 10. 6
6 1s22s22p6 K SHELL 1s2 2 e- L SHELL 2s2 2p6 8 e- PERIOD 2 GROUP VIIIA NOBLE GAS Ne
I KNOW AN ATOM WHOSE ATOMIC NUMBER IS 3. 7
7 1s22s1 K SHELL 1s2 2 e- L SHELL 2s1 1 e- PERIOD 2 GROUP IA METAL Li
GIVE THE ELECTRON CONFIGURATION OF THE FOLLOWING ELEMENTS, GIVE THE PERIOD NUMBER AND GROUP NUMBER TO WHERE IT BELONGS THEN DETERMINE WHETHER IT IS A METAL OR A NONMETAL
OXYGEN BERYLLIUM IODINE SULFUR BARIUM CARBON POTASSIUM O Be I S Ba C K 2, 6 2, 2 2, 8, 18, 18, 7 2, 8, 6 2, 8, 18, 18, 8, 2 2, 4 2, 8, 8, 1 P 2 G VIA P 2 G IIA P 5 G VIIA P 3 G VIA P 6 G IIA P 2 G IVA P 4 G IA

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BIOCHEMISTRY_3ELEMENTARY SUBSTANCES (2).ppt

  • 1.
    ELEMENTARY SUBSTANCES ELEMENTS AND THEPERIODIC TABLE OF ELEMENTS ATOMS: THE SMALLEST PARTICLE OF MATTER EVIDENCES OF SUBATOMIC PARTICLES DETERMINATION OF ATOMIC MASSES THE ELECTRONIC CONFIGURATION BY MAIN ENERGY LEVEL AND BY ENERGY SUBLEVEL PERIODICITY AND PERIODIC RELATIONSHIPS
  • 2.
    ELEMENT A PURE SUBSTANCEWHICH IS COMPOSED OF A SINGLE TYPE OF ATOM IT IS CONSISTS ONLY OF ATOMS THAT ALL HAVE THE SAME NUMBER OF PROTONS IN THEIR ATOMIC NUCLEI THE NUMBER OF PROTONS IN THE NUCLEUS IS THE DEFINING PROPERTY OF AN ELEMENT, AND IS REFERRED TO AS ITS ATOMIC NUMBER (z). ALL ATOMS WITH THE SAME ATOMIC NUMBER ARE ATOMS OF THE SAME ELEMENT IT CANNOT BE BROKEN DOWN INTO SIMPLER SUBSTANCES BY CHEMICAL MEANS. WHEN DIFFERENT ELEMENTS UNDERGO CHEMICAL REACTION, ATOMS ARE REARRANGED INTO NEW COMPOUNDS HELD TOGETHER BY CHEMICAL BONDS. 13p 14n 13e
  • 3.
    ELEMENTS CAN BEEXPRESSED USING CHEMICAL SYMBOLS. BORON GOLD TITANIUM SULFUR TANTALUM BARIUM URANIUM TITANIUM SULFUR TANTALUM
  • 4.
    EXERCISES. DECIPHER THEFOLLOWING USING CHEMICAL SYMBOLS SODIUM NEON THORIUM EINSTEINIUM PROTACTINI UM RHENIUM SULFUR PRASEODYM IUM INDIUM CERIUM SULFUR LANTHANUM URANIUM RHENIUM NITROGEN TELLURIUM MOLYBDENU M NICKEL COBALT MOLYBDENU M NITROGEN IODINE CARBON OXYGEN CARBON AMERICIUM POLONIUM IRIDIUM IODINE SULFUR NICKEL IODINE POTASSIUM POTASSIUM IODINE SULFUR AMERICIUM OXYGEN NITROGEN TELLURIUM PROTACTINI UM RADIUM IODINE SULFUR OXYGEN SODIUM, CARBON, IODINE, OXYGEN,
  • 5.
  • 6.
  • 7.
    METALLIC ELEMENTS Nearly allare shiny and grey-white in color Cu, Cs, and Au are shiny and golden Reflectivity is from intermediate to typically high. Almost all are solid in form. Rb, Cs, Fr, Cs and Hg are liquid at/near STP Densities are generally high. with some exceptions such as alkali metals Most metals are ductile and malleable. Some are brittle like Cr, Mn, Ge, Ru, W, Os, and Bi Most have hexagonal and cubic structure at freezing point. Close-packed crystal structures with high coordination numbers Atomic radius is from intermediate to very large. 112–298 pm (average 187 pm) Around half form allotropes Sn has a metalloid-like allotrope (grey Sn) which forms below 13.2 °C. Thermal conductivity is from medium to high Melting point is mostly high and generally expands during melting. Enthalpy of fusion is from low to high.
  • 8.
    s, p, d,and f periodic table block. Few number of electrons in outer s and p. except for Pd (0), Sn, Pb, and Fl (with 4), Bi (with 5) and Po (with 6) All has free electrons that facilitates electrical and thermal conductivity. Electrical conductivity is from good to high but falls gradually as temperature rises (as it turns to liquid) Over-all metallic behavior All metals tend to form cations. Metals seldom form covalent compounds (non-covalent bond) Oxidation number is nearly always positive Ionization energy is relatively low. Electronegativity is usually low. They are base-forming Their names usually end in –um or –ium with the following exceptions: Bi, W, Mn, Ni, and Co.
  • 9.
    nonMETALLI C ELEMENTS most are colourlessor dull red, yellow, green, or intermediate shades C, P, Se, and I are shiny and grey-white Reflectivity is from zero or low (mostly) to intermediate. most are gases. C, P, S, Se, and I are solid while Br is liquid. Their densities are often low. They are brittle when solid Some (C, P, S, and Se) have non-brittle forms. Their crystalline structure at freezing point ranges from hexagonal (H, He, C, N, and Xe), to cubic (O, F, Ne, P, Ar, Kr, Xe, and Rn) and orthorhombic ( S, Cl, Br & I). They have open structures & low coordination numbers. Atomic radius is from very small to intermediate 31–120 pm, average 76.4 pm
  • 10.
    s and p periodic tableblock. Most high number of electrons in the outer s and p except H and He. no, few, or directionally confined free electrons (generally hampering electrical and thermal conductivity) Electrical conductivity is from poor to good but increases as temperature rises Over-all nonmetallic behavior All nonmetals tend to form anions. Nonmetals form many covalent compounds Oxidation number is either positive or negative. Ionization energy and electronegativity are usually high. Some form allotropes while some (like graphitic C, black P and grey Se) are more metalloidal or metallic in nature Thermal conductivity is from almost negligible to very high. They are acid-forming Their names usually end in –on, -gen, or –ine with the following exceptions: S ends with –ur and Te and He end with -ium
  • 11.
    METALLOID S shiny and grey- whitein color Reflectivity is intermediate. All are solid in form Their densities are lower than nearby metals but higher than nearby nonmetals Their crystalline structure at freezing point ranges from rhombohedral (B, As, and Sb), to cubic (Si and Ge) and hexagonal (Te). relatively open crystal structures and medium coordination numbers[ Atomic radius is from small to intermediate: B, Si Ge, As, Sb, and Te 87–123 pm, average 115.5 pm They are brittle. all or nearly all form allotropes Some (red B and yellow As) are more nonmetallic in nature Si has high thermal conductivity but mostly are intermediate. Melting point is high and some contract, unlike (most) metals.
  • 12.
    p periodic table block. medium number(3–7) in the outer level valence electrons less freely delocalized; considerable covalent bonding present Electrical conductivity is from intermediate to good. Most behaves like metals. Over-all nonmetallic behavior Some have tendency to form anions in water. Solution chemistry dominated by formation and reactions of oxyanions. Form salts as well as covalent compounds Oxidation number is either positive or negative. Intermediate ionization energy and electronegativity is close to 2 (between 1.9 and 2.2)
  • 15.
    GIVE THE CHEMICALNAMES AND IDENTIFY WHETHER METAL, NONMETAL OR METALLOID 1. C 2. Ca 3. Cl 4. Ce 5. Cf 6. Cm 7. Co 8. Cr 9. Cs 10. Cu 11. S 12. Sb 13. Sc 14. Sm 15. Sn 16. Sr 17. Mg 18. Mn 19. Mo 20. Mt
  • 16.
  • 17.
    An element familyis a set of elements sharing common properties. Elements are classified into families because the three main categories of elements (metals, nonmetals and metalloids) are very broad. The characteristics of the elements in these families are determined primarily by the number of electrons in the outer energy shell. Element groups, on the other hand, are collections of elements categorized according to similar properties. Because element properties are largely determined by the behavior of valence electrons,
  • 18.
  • 19.
    PERIODS The rows ofthe table are known as periods. It is in the succesive periods that we observe the periodicity of properties of the elements. Each period has the full range of properties. For example more metallic elements occur to the left of a period, and the less metallic elements to the right; or oxides of the elements to the left are basic and acidic for elements to the right. The periods are simply numbered 1 though 7 from the top down GROUPS The columns of the table are known as groups or families. All the elements in a group have similar properties. Placing elements in groups is one of the most important ways of classifying them. There is some variation in properties within a group, but the changes are relatively small as one
  • 20.
    PERIOD 1 TO 7 GROUP 1TO 18 IA, IIA, IIIB, IVB, VB, VIB, VIIB, VIIIB, IIIA- VIIIA
  • 21.
  • 22.
    The periodic tablegroups are as follows (in the brackets are shown the old systems: European and American): Group 1 (IA): the alkali metals Group 2 (IIA): the alkaline earth metals Group 3 (IIIB) Group 4 (IVB) Group 5 (VB) Group 6 (VIB) Group 7 (VIIB) Group 8 (VIIIB) Group 9 (VIIIB) Group 10 (VIIIB) Group 11 (IB): the coinage metals (not a IUPAC-recommended name) Group 12 (IIB) Group 13 (IIIA): the boron group Group 14 (IVA): the carbon group Group 15 (VA): the pnictogens (not a IUPAC-recommended name) or nitrogen group Group 16 (VIA): the chalcogens
  • 23.
    ALKALI METALS arerecognized as a group and family of elements. These elements are metals. Hydrogen is not considered an alkali metal because the gas does not exhibit the typical properties of the group. However, under the right conditions of temperature and pressure, hydrogen can be an alkali metal. 1 valence electron Soft metallic solids Shiny, lustrous High thermal and electrical conductivity Low densities, increasing with atomic mass Relatively low melting points, decreasing with atomic mass Vigorous exothermic reaction with water to GROU P IA (CAS) GROU P 1 (IUPAC)
  • 24.
    ALKALINE EARTH METALSor simply alkaline earths are recognized as an important group and family of elements. These elements are metals. Examples include calcium and magnesium. 2 valence electrons Metallic solids, harder than the alkali metals Shiny, lustrous, oxidize easily High thermal and electrical conductivity More dense than the alkali metals Higher melting points than alkali metals Exothermic reaction with water, increasing as you move down the group; beryllium does not react with water; magnesium reacts only with steam GROUP IIA (CAS) GROUP 2 (IUPAC)
  • 25.
    TRANSITION METALS isthe largest family of elements. The center of the periodic table contains the transition metals, plus the two rows below the body of the table (lanthanides and actinides) are special transition metals. Groups 3-12 The d and f block metals have 2 valence electrons Hard metallic solids Shiny, lustrous High thermal and electrical conductivity Dense High melting points Large atoms exhibit GROUP IIIB GROUP 3 GROUP VIIIB GROUP 8, 9, 10 GROUP IB / IIB GROUP 11, 12
  • 26.
    BORON GROUP orearth metal family is not as well-known as some of the other element families. 3 valence electrons Diverse properties, intermediate between those of metals and nonmetals CARBON GROUP is made up of elements called tetrels, which refers to their ability to carry a charge of 4. 4 valence electrons Diverse properties, intermediate between those of metals and nonmetals GROUP IIIA GROUP 13 GROUP IVA GROUP 14
  • 27.
    NITROGEN GROUP orpnictogens is a significant element family. 5 valence electrons Diverse properties, intermediate between those of metals and nonmetals OXYGEN GROUP or the chalcogens family. 6 valence electrons Diverse properties, changing from nonmetallic to metallic in moving down the family GROUP VA (CAS) GROUP 15 (IUPAC) GROUP VIA (CAS) GROUP
  • 28.
    HALOGEN FAMILY isa group of reactive nonmetals. 7 valence electrons Reactive nonmetals Melting points and boiling points increase with increasing atomic number High electron affinities Change state as it moves down the family, with fluorine and chlorine existing as gases at room temperature while bromine is a liquid and iodine is a solid NOBLE GASES are a family of nonreactive nonmetals. 8 valence electrons Typically exist as monatomic gases, although these elements do (rarely) form compounds Stable electron octet makes nonreactive (inert) under GROUP VIIA (CAS) GROUP 17 (IUPAC) GROUP VIIIA (CAS) GROUP 18 (IUPAC)
  • 29.
    ATOMS • incredibly tiny •numerous • perpetually in motion • ageless • An atom takes part in chemical reactions independently. • An atom can be divided into a number of sub-atomic particles - electron, proton and neutron.
  • 30.
  • 31.
    • ELECTRONS – discoveredby Sir John Joseph Thomson in 1897 – are located in an electron cloud, which is the area surrounding the nucleus of the atom – can abbreviated as e- – have a negative charge that is equal in magnitude to the positive charge of the protons – mass is considerably less than that of a proton or neutron
  • 32.
    • PROTONS – discoveredby Ernest Rutherford in the year 1919, when he performed his gold foil experiment. – exist in a nucleus and have a positive nuclear charge. – atomic number or proton number is the number of protons present in an atom – The atomic number determines an element (e.g., the element of atomic number 6 is carbon).
  • 33.
    • NEUTRONS – werediscovered by James Chadwick in 1932, when he demonstrated that penetrating radiation incorporated beams of neutral particles – located in the nucleus with the protons – Along with protons, they make up almost all of the mass of the atom – The number of neutrons is called the neutron number and can be found by subtracting the proton number from the atomic mass number.
  • 34.
    • u isthe SI symbol for atomic mass unit. • The positive charge of protons cancels the negative charge of the electrons. Neutrons have no charge. • With regard to mass, protons and neutrons are very similar, and have a much greater mass than electrons. Compared with neutrons and protons, the mass of an electron is usually negligible. • Spin is associated with the rotation of a particle. Protons, neutrons, and electrons each have a total spin of 1/2.
  • 35.
    RELATIONSHIP OF THEATOMIC MASS AND ATOMIC NUMBER WITH THE SUBATOMIC PARTICLES • A = p + n - Atomic mass (A) is the sum of the number of protons (p) and neutrons (n) in the nucleus of an atom. • Z = p - Atomic number (Z) gives the number of protons (p) in the nucleus. It is what distinguishes the atoms of a given element. • n = A – p - The number of neutrons can be found by subtracting the proton number from the atomic mass number.
  • 36.
    Element Atomic Mass (A) Atomic Number (Z) Protons (p) Electrons ( e-) Neutrons (n) K 39 19 19 19 20 Ar 40 18 18 18 22 Na 23 11 11 11 12 O 16 8 8 8 8 Au 197 79 Ag 108 47 Zn Mg Og Mn I 79 79 118 47 47 61
  • 37.
    ISOTOPES: Not allatoms are alike • Isotopes are atoms of the same element with different number of neutrons or different atomic masses • Isotopes are atoms with the same number of neutrons. Isotopes are different forms of a single element.
  • 38.
    Known Isotopes ofan Element • Hydrogen - 1H - 2H - 3H - 4H - 5H -6H -7H  Lithium - 3Li - 8Li - 4Li - 9Li - 5Li - 10Li - 6Li - 11Li - 7Li - 12Li  Beryllium - 5Be - 12Be - 6Be - 13Be - 7Be - 14Be - 8Be - 15Be - 9Be - 16Be - 10Be - 11Be  Boron - 6B - 12B - 7B - 13B - 8B - 14B - 9B - 15B - 10B - 16B - 11B - 17B  Carbon - 8C - 14C - 9C - 15C - 10C - 16C - 11C - 17C - 12C - 18C - 13C - 19C particle p e- n 1H 1 1 0 2H 1 1 1 3H 1 1 2 particle p e- n 3Li 3 3 0 4Li 3 3 1 5Li 3 3 2 particle p e- n 5Be 4 4 1 6Be 4 4 2 7Be 4 4 3
  • 39.
    HOW IS ATOMICMASS DETERMINED? A = (IM X FA)1 + (IM X FA)2 + (IM X FA)3 + ....... IM is the isotopic mass of a naturally occuring element FA is the isotopic abundance of a naturally occuring element What is the average atomic mass of magnesium? ISOTOPE ISOTOPIC ABUNDANCE ISOTOPIC MASS Mg - 24 78.99 % 23.98501417 u Mg - 25 10.00 % 24.98583692 u Mg - 26 11.01 % 25.98259292 u A = (IM X FA)Mg-24 + (IM X FA)Mg-25 + (IM X FA)Mg-26 A = (23.98501417 X 0.7899) + (24.98583692 X 0.1000) + (25.98259292 X 0.1101) A = 24.30502986 A = 24.31 u
  • 40.
    Calculate the atomicmass of potassium given the following: K-39 isotopic abundance- 93.12% isotopic mass- 38.964 u K-41 isotopic abundance- 6.88% isotopic mass- 40.962 u A = (IM X FA)K-39 + (IM X FA)K-41 A = (38.964 X 0. 9312) + (40.962 X 0.0688) A = 39.102 u A = 39.1 u Use the atomic masses of each of the two isotopes of chlorine along with their isotopic abundances to calculate the average atomic mass of chlorine. chlorine-35: atomic mass = 34.969 amu and % abundance = 75.77% chlorine-37: atomic mass = 36.966 amu and % abundance = 24.23% A = (IM X FA)Cl-35 + (IM X FA)Cl-37 A = (34.969 X 0.7577) + (36.966 X 0.2423) A = 35.457 u A = 35.46 u
  • 41.
    ELEMENTAL IONS: chargedparticles • ELEMENTAL IONS ARE PARTICLES THAT RESULTED FROM THE LOSS OR GAIN OF ELECTRON. • WHEN AN ATOM LOSES ELECTRON IT WILL BECOME POSITIVELY CHARGED (CATION). • Na has 11p, 11e-, and 12n. When one of its electron is transferred to another atom, it will become positively charged. The number of p and n are the same but the number of electron becomes 10. The charge of the this ion can be determined by subtracting the number of electrons from the number of proton. Charge = p – e- = 11 – 10 = 1 • Na atom  Na+ ion
  • 42.
    • WHEN ANATOM GAINS ELECTRON IT WILL BECOME NEGATIVELY CHARGED (ANION). • Cl has 17p, 17e-, and 18n. When one electron is accepted by Cl , it will become negatively charged. The number of p and n are the same but the number of electron becomes 18. The charge of the this ion can be determined by subtracting the number of electrons from the number of proton. • Charge = p – e- = 17 – 18 = -1 • Cl atom  Cl- ion
  • 43.
    PARTICLE p e-n F fluorine atom 9 9 10 O oxygen atom 8 8 8 N nitrogen atom 7 7 7 S sulfur atom 16 16 16 PARTICLE p e- n F- fluoride ion 9 10 10 O2- oxide ion 8 10 8 N3- nitride ion 7 10 7 S2- sulfide ion 16 18 16
  • 44.
    ELECTRON CONFIGUR ATION The electron configurationof an atomic species (neutral or ionic) allows us to understand the shape and energy of its electrons. Many general rules are taken into consideration when assigning the location of the electron to its prospective energy state, however these assignments are arbitrary and it is always uncertain as to which electron is being described. Knowing the electron
  • 45.
  • 46.
    The protons andneutrons are found inside the nucleus of an atom, while the electrons are constantly moving around the nucleus. The nucleus and the electrons interact to form the most stable arrangement possible, the ground state, which requires the least amount of energy. There are four main concepts that must remembered in electron configuration. These are the energy levels, the enrgy sublevel, the orbitals, and number of electrons per sublevel.
  • 47.
    Energy Levels andthe Atomic Model The concept of energy levels is one part of the atomic model that is based on a mathematical analysis of atomic spectra. Each electron in an atom has an energy signature that is determined by its relationship with other negatively charged electrons in the atom and the positively charged atomic nucleus. An electron can change energy levels, but only by steps or quanta, not continuous increments. The energy of an energy level increases the further out from the nucleus it is. The lower the number of a principal energy level, the closer together the electrons are to each other and to the nucleus of the atom. During chemical reactions, it's more difficult to remove an electron from a lower energy level
  • 48.
    Rules of PrincipalEnergy Levels A principal energy level may contain up to 2n2 electrons, with n being the number of each level. The first energy level can contain 2(1)2 or two electrons; the second can contain up to 2(2)2 or eight electrons; the third can contain up to 2(3)2 or 18 electrons, and so on. PRINCIPAL ENERGY LEVEL. n SHELL LETTER MAXIMUM NUMBER OF ELECTRONS, 2n2 1 K 2n2 = 2(1)2 = 2(1) = 2 2 L 2n2 = 2(2)2 = 2(4) = 8 3 M 2n2 = 2(3)2 = 2(9) = 18 4 N 2n2 = 2(4)2 = 2(16) = 32 5 O 2n2 = 2(5)2 = 2(25) = 50 6 P 2n2 = 2(6)2 = 2(36) = 72 7 Q 2n2 = 2(7)2 = 2(49) = 98
  • 49.
    ENERGY SUBLEVELS/ENERGY SUBSHELLS Each energy level,or shell, is divided into sublevels. The terms sublevel and subshell are used interchangeably. The sublevels are represented by the letters s, p, d, and f. Each energy level has certain sublevels. ENERGY LEVELS ENERGY SUBLEVELS n = 1 s n = 2 s, p n = 3 s, p, d n = 4 s, p, d, f Energy levels that are higher than four would contain additional sublevels such as g and h. No known atom, in its ground state, would have electrons that
  • 50.
    ORBITALS Each sublevel ismade up of orbitals. Each sublevel has a different number of orbitals. ENERGY SUBLEVELS ORBITALS s 1 p 3 d 5 f 7 Orbitals are classified by their shape. s orbitals are spherical in shape.
  • 51.
    Orbitals found inthe p subshell are dumbbell, or figure-eight, shaped. Remember, in each p subshell there are three orbitals, so there are three dumbbell shapes. Atoms are three dimensional objects, and the orbitals are oriented around the different axes. The image below shows the three p orbitals and the complete p sublevel containing all three orbitals.
  • 52.
    The first principalenergy level has one sublevel that contains one orbital, called the s orbital. The s orbital can contain a maximum of two electrons. The next principal energy level contains one s orbital and three p orbitals. The set of three p orbitals can hold up to six electrons. Thus, the second principal energy level can hold up to eight electrons, two in the s orbital and six in the p orbital. The third principal energy level has one s orbital, three p orbitals, and five d orbitals, which can each hold up to 10 electrons. This allows for a maximum of 18 electrons. The fourth and higher levels have an f sublevel in addition to the s, p, and d orbitals. The f sublevel contains seven f orbitals, which can each hold up to 14 electrons. The total number of electrons in
  • 53.
    PRINCIPAL ENERGY LEVEL, n NUMBEROF ORBITALS ORBITALS NUMBER OF ELECTRONS PER ORBITAL 1 1 s 2 2 2 s p 2 6 3 3 s p d 2 6 10 4 4 s p d f 2 6 10 14 Electron Notation The notation used to indicate the type of energy level and the number of electrons in that level has a coefficient for the number of the principal energy level, a letter for the sublevel, and a superscript for the number of electrons located in that sublevel. For example, the notation 4p3 indicates the fourth principal energy level, the p sublevel, and the presence of
  • 54.
    An orbital diagramhelps to determine the electron configuration of an element. An element’s electron configuration is the arrangement of the electrons in the shells. There are a few guidelines for working out this configuration: 1. Each orbital can hold only two electrons. Electrons that occur together in an orbital are called an electron pair. 2. An electron will always try to enter the orbital with the lowest energy. 3. An electron can occupy an orbital on its own, but it would rather occupy a lower-energy orbital with another electron before occupying a higher- energy orbital. In other words, within one energy
  • 55.
    THE BUILDING-UP PRINCIPLE The Aufbauprinciple determines an atom’s electron configuration by adding electrons to atomic orbitals following a defined set of rules. Although the nucleus of an atom is very dense, the electrons around it can take on a variety of positions which can be summarized as an electron configuration. An element’s electron configuration can be represented using energy level diagrams, or Aufbau diagrams. The Aufbau principle (from the German Aufbau meaning “building up, construction”) describes a model-building method in which an atom is “built up” by progressively
  • 56.
    Filling in anAufbau Diagram The order in which orbitals are filled is given by the Madelung rule. The rule is based on the total number of nodes in the atomic orbital, n + ℓ, which is related to the energy. In this context, n represents the principal quantum number and ℓ represents the azimuthal quantum number. The values ℓ = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. According to the principle, electrons fill orbitals starting at the lowest available energy states before filling higher states (e.g., 1s before 2s). The Madelung rule defines the order in which atomic orbitals are filled with electrons. Electrons fill orbitals starting at the lowest available energy state before filling higher states.
  • 57.
  • 58.
    The following stepsdetail how to draw an Aufbau diagram 1. Determine the number of electrons that the atom has. 2. Fill the s orbital in the first energy level (the 1s orbital) with the first two electrons. 2. Fill the s orbital in the second energy level (the 2s orbital) with the second two electrons. 3. Put one electron in each of the three p orbitals in the second energy level (the 2p orbitals) and then if there are still electrons remaining, go back and place a second
  • 59.
    Limitations to Aufbau TheAufbau principle is based on the idea that the order of orbital energies is fixed— both for a given element and between different elements. This assumption is approximately true—enough for the principle to be useful—but not physically reasonable. It models atomic orbitals as “boxes” of fixed energy into which at most two electrons can be placed. However, the energy of an electron in an atomic orbital depends on the energies of all the other electrons of the atom.
  • 60.
    Electron Configuration Standard Notation Aspecial type of notation is used to write an atom’s electron configuration. The notation describes the energy levels, orbitals, and the number of electrons in each. For example, the electron configuration of lithium is 1s22s1. The number and letter describe the energy level and orbital, and the number above the orbital shows how many electrons are in that orbital. Using standard notation, the electron configuration of fluorine is 1s22s22p5.
  • 61.
    Electrons will fillthe lowest energy orbitals first and then move up to higher energy orbitals only after the lower energy orbitals are full.. Although the implications are clear for orbitals of different principal quantum number (n), which are clearly of different energy, the filling order is less clear for degenerate sublevels. For example, for boron through neon, the electron filling order of the 2p orbitals follows Hund’s Rule. Hund’s Rule states that: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied. All of the electrons in singly occupied orbitals have the same spin.
  • 63.
    Diamagnetism Any time twoelectrons share the same orbital, their spin quantum numbers have to be different. In other words, one of the electrons has to be “spin- up,”, while the other electron is “spin-down,” This is important when it comes to determining the total spin in an electron orbital. In order to decide whether electron spins cancel, add their spin quantum numbers together. Whenever two electrons are paired together in an orbital, or their total spin is 0, they are called diamagnetic electrons.
  • 64.
    Paramagnetism Electrons that arealone in an orbital are called paramagnetic electrons. Remember that if an electron is alone in an orbital, the orbital has a net spin, because the spin of the lone electron does not get canceled out. If even one orbital has a net spin, the entire atom will have a net spin. Therefore, an atom is considered to be paramagnetic when it contains at least one paramagnetic electron. In other words, an atom could have 10 paired (diamagnetic) electrons,
  • 65.
    - Periodic trendsare specific patterns that are present in the periodic table that illustrate different aspects of a certain element, including its size and its electronic properties. - Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties. - These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and Periodic Trends
  • 66.
    property describing anatom's ability to attract and bind with electrons. Because electronegativity is a qualitative property, there is no standardized method for calculating electronegativity. However, the most common scale for quantifying electronegativity is the Pauling scale, named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature of electronegativity. Electronegativity values for each
  • 67.
    From left toright across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. If the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one. From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic
  • 68.
    Important exceptions ofthe above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values. As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.
  • 69.
    Ionization Energy - isthe energy required to remove an electron from a neutral atom in its gaseous phase. Conceptually, ionization energy is the opposite of electronegativity. The lower this energy is, the more readily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely it is that the atom becomes a cation. Generally, elements on the right side of the periodic table have a higher ionization energy because their valence shell is nearly filled. Elements on the left side of the periodic table have low ionization
  • 70.
    The ionization energyof the elements within a period generally increases from left to right. This is due to valence shell stability. The ionization energy of the elements within a group generally decreases from top to bottom. This is due to electron shielding. The noble gases possess very high ionization energies because of their full valence shells as indicated in the graph. Note that helium has the highest ionization energy of all the elements.
  • 71.
    Electron Affinity - isthe ability of an atom to accept an electron. Unlike electronegativity, electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom's affinity for electrons. - Electron affinity increases from left to right within a period. This is caused by the decrease in atomic radius. - Electron affinity decreases from top to bottom within a group. This is caused by the increase in
  • 72.
    ATOMIC RADIUS - Whilethe atomic radius can be defined in a number of different ways, the general atomic radius trend across the periodic table holds true. The atomic radius for atoms of an element tends to go up as you move down a group of elements in the table. The atomic radius increases as you move down a column because for every new row of the table a new electron shell is added to the atom. - Atomic radius decreases from left to right within a period. This is caused by the increase in the number of protons and electrons across a period. One proton has a greater effect than one electron; thus, electrons are pulled towards the nucleus, resulting in a smaller radius. Atomic radius increases from top to bottom within a
  • 73.
    Melting Point - isthe amount of energy required to break a bond(s) to change the solid phase of a substance to a liquid. Generally, the stronger the bond between the atoms of an element, the more energy required to break that bond. Because temperature is directly proportional to energy, a high bond dissociation energy correlates to a high temperature. Melting points are varied and do not generally form a distinguishable trend across the periodic table. Metals generally possess a high melting point. Most non-metals possess low melting points. The non-metal carbon possesses the highest melting point of all the elements. The semi-metal boron also possesses a high melting point.
  • 74.
    The Metallic Character -of an element can be defined as how readily an atom can lose an electron. From right to left across a period, metallic character increases because the attraction between valence electrons and the nucleus is weaker, enabling an easier loss of electrons. Metallic character increases as you move down a group because the atomic size is increasing. When the atomic size increases, the outer shells are farther away. The principal quantum number increases and average electron density moves farther
  • 75.
    from left toright across a period. This is caused by the decrease in radius of the atom that allows the outer electrons to ionize more readily. Metallic characteristics increase down a group. Electron shielding causes the atomic radius to increase thus the outer electrons ionizes more readily than electrons in smaller atoms. Metallic character relates to the ability to lose electrons, and nonmetallic character relates to the
  • 77.
  • 78.
    ELECTRON CONFIGURATION BY MAINENERGY LEVEL THE DISTRIBUTION OR ARRANGEMENT OF ELECTRONS IN THE ATOM’S SHELL. ELECTRONS ARE DISTRIBUTED STARTING FROM THE INNERMOST SHELL. THERE ARE SEVEN SHELLS OR MAIN ENERGY LEVELS WHICH ARE REPRESENTED BY EITHER SHELL LETTER OR SHELL NUMBER.
  • 79.
    I KNOW ANATOM WHOSE ATOMIC NUMBER IS 38. 1
  • 80.
    1 1s22s22p63s23p64s23d104p65s2 K SHELL 1s2 2e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p6 3d10 18 e- N SHELL 4s2 4p6 8 e- O SHELL 5s2 2 e- PERIOD 5 GROUP IIA METAL Sr
  • 81.
    Psych e- Fills MEL IKNOW AN ATOM WHOSE ATOMIC NUMBER IS 33. 2
  • 82.
    2 1s22s22p63s23p64s23d104p3 K SHELL 1s2 2e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p6 3d10 18 e- N SHELL 4s2 4p3 5 e- PERIOD 4 GROUP VA NONMETAL As
  • 83.
    I KNOW ANATOM WHOSE ATOMIC NUMBER IS 20. 3
  • 84.
    3 1s22s22p63s23p64s2 K SHELL 1s2 2e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p6 8 e- N SHELL 4s2 2 e- PERIOD 4 GROUP IIA METAL Ca
  • 85.
    I KNOW ANATOM WHOSE ATOMIC NUMBER IS 17. 4
  • 86.
    4 1s22s22p63s23p5 K SHELL 1s2 2e- L SHELL 2s2 2p6 8 e- M SHELL 3s2 3p5 7 e- PERIOD 3 GROUP VIIA NONMETAL Cl
  • 87.
    I KNOW ANATOM WHOSE ATOMIC NUMBER IS 11. 5
  • 88.
    5 1s22s22p63s1 K SHELL 1s2 2e- L SHELL 2s2 2p6 8 e- M SHELL 3s1 1 e- PERIOD 3 GROUP IA METAL Na
  • 89.
    I KNOW ANATOM WHOSE ATOMIC NUMBER IS 10. 6
  • 90.
    6 1s22s22p6 K SHELL 1s2 2e- L SHELL 2s2 2p6 8 e- PERIOD 2 GROUP VIIIA NOBLE GAS Ne
  • 91.
    I KNOW ANATOM WHOSE ATOMIC NUMBER IS 3. 7
  • 92.
    7 1s22s1 K SHELL 1s2 2e- L SHELL 2s1 1 e- PERIOD 2 GROUP IA METAL Li
  • 93.
    GIVE THE ELECTRON CONFIGURATIONOF THE FOLLOWING ELEMENTS, GIVE THE PERIOD NUMBER AND GROUP NUMBER TO WHERE IT BELONGS THEN DETERMINE WHETHER IT IS A METAL OR A NONMETAL
  • 94.
    OXYGEN BERYLLIUM IODINE SULFUR BARIUM CARBON POTASSIUM O Be I S Ba C K 2, 6 2, 2 2,8, 18, 18, 7 2, 8, 6 2, 8, 18, 18, 8, 2 2, 4 2, 8, 8, 1 P 2 G VIA P 2 G IIA P 5 G VIIA P 3 G VIA P 6 G IIA P 2 G IVA P 4 G IA