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Ch 5 Notes

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  • 1. Elements, Atoms, and Ions
  • 2.  
  • 3. The Language of Chemistry
    • CHEMICAL ELEMENTS -
      • pure substances that cannot be decomposed by ordinary means to other substances.
    Sodium Bromine Aluminum
  • 4. The Language of Chemistry
    • The elements, their names, and symbols are given on the PERIODIC TABLE
    • How many elements are there?
  • 5. The Periodic Table
    • Dmitri Mendeleev (1834 - 1907)
  • 6. Glenn Seaborg (1912-1999 )
    • Discovered 8 new elements.
    • Only living person for whom an element was named.
  • 7.
    • An atom consists of a
    • nucleus
      • (of protons and neutrons )
    • electrons in space about the nucleus.
    The Atom http:// www.youtube.com/watch?v =kypne21A0R4 http:// www.youtube.com/watch?v =qmgE0w6E6ZI&mode= related&search = Nucleus Electron cloud
  • 8. Copper atoms on silica surface.
    • An atom is the smallest particle of an element that has the chemical properties of the element.
    Distance across = 1.8 nanometer (1.8 x 10 -9 m)
  • 9. Subatomic Particles
    • Quarks
      • component of protons & neutrons
      • 6 types
      • 3 quarks = 1 proton or 1 neutron
    He
  • 10. The red compound is composed of • nickel (Ni) (silver) • carbon (C) (black) • hydrogen (H) (white) • oxygen (O) (red) • nitrogen (N) (blue)
    • CHEMICAL COMPOUNDS are composed of atoms and so can be decomposed to those atoms.
  • 11. Compounds
      • composed of 2 or more elements in a fixed ratio
      • properties differ from those of individual elements
      • EX : table salt (NaCl)
  • 12. A MOLECULE is the smallest unit of a compound that retains the chemical characteristics of the compound.
    • Composition of molecules is given by a MOLECULAR FORMULA
    H 2 O C 8 H 10 N 4 O 2 - caffeine
  • 13. ELEMENTS THAT EXIST AS DIATOMIC MOLECULES Remember: BrINClHOF These elements only exist as PAIRS. Note that when they combine to make compounds, they are no longer elements so they are no longer in pairs!
  • 14. Where Did it All Begin?
    • The word “atom” comes from the Greek word “atomos” which means indivisible .
    • The idea that all matter is made up of atoms was first proposed by the Greek philosopher Democritus in the 5th century B.C.
  • 15. Dalton’s Atomic Theory
    • John Dalton (1766-1844) proposed an atomic theory
    • While this theory was not completely correct, it revolutionized how chemists looked at matter and brought about chemistry as we know it today instead of alchemy
    • Thus, it’s an important landmark in the history of science.
  • 16. Dalton’s Atomic Theory - Summary
    • matter is composed, indivisible particles (atoms)
    • all atoms of a particular element are identical
    • different elements have different atoms
    • atoms combine in certain whole-number ratios
    • In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.
  • 17. Problems with Dalton’s Atomic Theory?
    • 1. matter is composed, indivisible particles
    • Atoms Can Be Divided, but only in a nuclear reaction
    • 2. all atoms of a particular element are identical
    • Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)!
    • 3. different elements have different atoms
    • YES!
    • 4. atoms combine in certain whole-number ratios
    • YES! Called the Law of Definite Proportions
    • 5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.
    • Yes, except for nuclear reactions that can change atoms of one element to a different element
  • 18.
    • J.J. Thompson discovered electrons in 1897 by performing the cathode ray experiment.
    • Plum Pudding Model
    Cathode Ray tube What about Electrons?
    • Much of Dalton’s theories are accepted except that atoms are known now to be divisible into sub-particles.
  • 19.
    • Robert A. Millikin 1868-1953 carried out experiments to discover the charge and mass of electrons by
    • Physicist James Chadwick 1891-1974 confirmed the existence of yet another subatomic particle: the neutron
    Millikin Oil Drop experiment
  • 20. Rutherford’s experiment. Ernest Rutherford wondered how all these particles were put together
  • 21.
    • The modern view of the atom was developed by Ernest Rutherford (1871-1937).
  • 22. Results of foil experiment if Plum Pudding model had been correct .
  • 23. What Actually Happened
  • 24. Niels Bohr
    • Niels Bohr applies quantum theory to Rutherford's atomic structure by assuming that electrons travel in stationary orbits defined by their angular momentum.
    • This led to the calculation of possible energy levels for these orbits and the postulation that the emission of light occurs when an electron moves into a lower energy orbit.
  • 25. Niels Bohr
  • 26. ATOM COMPOSITION
    • protons and neutrons in the nucleus.
    • the number of electrons is equal to the number of protons.
    • electrons in space around the nucleus.
    • extremely small. One teaspoon of water has 3 times as many atoms as the Atlantic Ocean has teaspoons of water.
    The atom is mostly empty space
  • 27. ATOMIC COMPOSITION
    • Protons (p + )
      • + electrical charge
      • mass = 1.672623 x 10 -24 g
      • relative mass = 1.007 atomic mass units (amu) but we can round to 1
    • Electrons (e - )
      • negative electrical charge
      • relative mass = 0.0005 amu but we can round to 0
    • Neutrons (n o )
      • no electrical charge
      • mass = 1.009 amu but we can round to 1
  • 28. Atomic Number, Z
    • All atoms of the same element have the same number of protons in the nucleus, Z
    13 Al 26.981 Atomic number Atom symbol AVERAGE Atomic Mass
  • 29. Mass Number, A
    • C atom with 6 protons and 6 neutrons is the mass standard
    • = 12 atomic mass units
    • Mass Number (A) = # protons + # neutrons
    • NOT on the periodic table…(it is the AVERAGE atomic mass on the table)
    • A boron atom can have A = 5 p + 5 n = 10 amu
  • 30. Isotopes
    • Atoms of the same element (same Z) but different mass number (A).
    • Boron-10 ( 10 B) has 5 p and 5 n
    • Boron-11 ( 11 B) has 5 p and 6 n
    10 B 11 B
  • 31. Figure 3.10: Two isotopes of sodium.
  • 32. Isotopes & Their Uses Bone scans with radioactive technetium-99.
  • 33. Isotopes & Their Uses The tritium content of ground water is used to discover the source of the water, for example, in municipal water or the source of the steam from a volcano.
  • 34. Atomic Symbols
    • Show the name of the element, a hyphen, and the mass number in hyphen notation
    • sodium-23
    • Show the mass number and atomic number in nuclear symbol form
    • mass number
    • 23 Na
    • atomic number 11
  • 35. Isotopes?
    • Which of the following represent isotopes of the same element? Which element?
    • 234 X 234 X 235 X 238 X
    • 92 93 92 92
  • 36. Answer:
    • 234 U 234 Np 235 U 238 U
    • 92 93 92 92
    • 234 Np is not an isotope of Uranium.
    • 93
  • 37. Counting Protons, Neutrons, and Electrons
    • Protons: Atomic Number (from periodic table)
    • Neutrons: Mass Number minus the number of protons (mass number is protons and neutrons because the mass of electrons is negligible)
    • Electrons:
      • If it’s an atom, the protons and electrons must be the SAME so that it is has a net charge of zero (equal numbers of + and -)
      • If it does NOT have an equal number of electrons, it is not an atom, it is an ION. For each negative charge, add an extra electron. For each positive charge, subtract an electron (Don’t add a proton!!! That changes the element!)
  • 38. Learning Check – Counting
    • Naturally occurring carbon consists of three isotopes, 12 C, 13 C, and 14 C. State the number of protons, neutrons, and electrons in each of these carbon atoms.
    • 12 C 13 C 14 C
    • 6 6 6
    • #p + _______ _______ _______
    • #n o _______ _______ _______
    • #e - _______ _______ _______
  • 39. Answers
    • 12 C 13 C 14 C
    • 6 6 6
    • #p + 6 6 6
    • #n o 6 7 8
    • #e - 6 6 6
  • 40. Learning Check
    • An atom has 14 protons and 20 neutrons.
    • A. Its atomic number is
    • 1) 14 2) 16 3) 34
    • B. Its mass number is
    • 1) 14 2) 16 3) 34
    • C. The element is
    • 1) Si 2) Ca 3) Se
    • D. Another isotope of this element is
    • 1) 34 X 2) 34 X 3) 36 X
    • 16 14 14
  • 41. Solution
    • An atom has 14 protons and 20 neutrons.
    • A. It has atomic number
    • 1) 14
    • B. It has a mass number of
    • 3) 34
    • C. The element is
    • 1) Si
    • D. Another isotope of this element would be
    • 3) 36 X
    • 14
  • 42. IONS
    • IONS are atoms or groups of atoms with a positive or negative charge.
    • Taking away an electron from an atom gives a CATION with a positive charge
    • Adding an electron to an atom gives an ANION with a negative charge .
    • To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! Examples: Na + Ca +2 I - O -2
    • Na Ca I O
  • 43. Forming Cations & Anions A CATION forms when an atom loses one or more electrons. An ANION forms when an atom gains one or more electrons Mg --> Mg 2+ + 2 e- F + e- --> F -
  • 44. PREDICTING ION CHARGES
    • In general
    • metals (Mg) lose electrons ---> cations
    • nonmetals (F) gain electrons ---> anions
  • 45. Charges on Common Ions +3 By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom. http:// www.youtube.com/watch?v = IpaEGhjpZgc -4 -1 -2 -3 +1 +2
  • 46. Learning Check – Counting
    • State the number of protons, neutrons, and electrons in each of these ions.
    • 39 K + 16 O -2 41 Ca +2
    • 19 8 20
    • #p + ______ ______ _______
    • #n o ______ ______ _______
    • #e - ______ ______ _______
  • 47. Learning Check – Counting
    • State the number of protons, neutrons, and electrons in each of these ions.
    • 39 K + 16 O -2 41 Ca +2
    • 19 8 20
    • #P 19 8 20
    • #N 20 8 21
    • #E 18 10 18
  • 48. One Last Learning Check
    • Write the nuclear symbol form for the following atoms or ions:
    • A. 8 p + , 8 n, 8 e - ___________
    • B. 17p + , 20n, 17e - ___________
    • C. 47p + , 60 n, 46 e - ___________
  • 49. Answers
    • Write the nuclear symbol form for the following atoms or ions:
    • A. 8 p + , 8 n, 8 e - 16 O
    • 8
    • B. 17p + , 20n, 17e - 37 Cl
    • 17
    • C. 47p + , 60 n, 46 e - 107 Ag +
    • 47
  • 50. AVERAGE ATOMIC MASS
    • Because of the existence of isotopes, the mass of a collection of atoms has an average value.
    • Boron is 20% 10 B and 80% 11 B. That is, 11 B is 80 percent abundant on earth.
    • For boron atomic weight
    • = 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu
    10 B 11 B
  • 51. Isotopes & Average Atomic Mass
    • Because of the existence of isotopes, the mass of a collection of atoms has an average value.
    • 6 Li = 7.5% abundant and 7 Li = 92.5%
      • Avg. Atomic mass of Li = ______________
    • 28 Si = 92.23%, 29 Si = 4.67%, 30 Si = 3.10%
      • Avg. Atomic mass of Si = ______________
  • 52. The Periodic Table
  • 53. Periods in the Periodic Table
  • 54. Groups in the Periodic Table Elements in groups react in similar ways!
  • 55. Regions of the Periodic Table
  • 56. University of North Texas Chemistry Building The only full periodic table with all the elements (except man-made) in existence
  • 57. Group 1A: Alkali Metals Cutting sodium metal Reaction of potassium + H 2 O
  • 58. Magnesium Magnesium oxide Group 2A: Alkaline Earth Metals
  • 59. Group 7A: The Halogens (salt makers) F, Cl, Br, I, At
  • 60. Group 8A: The Noble (Inert) Gases He, Ne, Ar, Kr, Xe, Rn
    • Lighter than air balloons
    • “ Neon” signs
    • Very Unreactive because they have full electron levels
    XeOF 4
  • 61. Transition Elements
    • Lanthanides and actinides
    Iron in air gives iron(III) oxide
  • 62.
    • But Why Does The Periodic Table end too quickly? Can’t we just add a proton to the end???
    • Let’s Watch!
  • 63. Chemistry of Groups
    • Group 1: Alkali Metals
    • Group 2: Alkaline Earth Metals
    • Groups 3-11: Transition Elements
    • Group 17: Halogens
    • Group 18: Noble Gases
    • Diatomic Molecules
  • 64. Group 1: Alkali Metals
    • Most active metals, only found in compounds in nature
    • React violently with water to form hydrogen gas and a strong base: 2 Na (s) + H 2 O (l)  2 NaOH (aq) + H 2 (g)
    • 1 valence electron
    • Form +1 ion by losing that valence electron
    • Form oxides like Na 2 O, Li 2 O, K 2 O
  • 65. Alkali Metals
    • The alkali metals, found in group 1 of the periodic table (formerly known as group IA), are very reactive metals that do not occur freely in nature. These metals have only one electron in their outer shell. Therefore, they are ready to lose that one electron in ionic bonding with other elements.
    • As with all metals, the alkali metals are malleable, ductile, and are good conductors of heat and electricity. The alkali metals are softer than most other metals.
    • Cesium and francium are the most reactive elements in this group. Alkali metals can explode if they are exposed to water.
  • 66. Group 2: Alkaline Earth Metals
    • Very active metals, only found in compounds in nature
    • React strongly with water to form hydrogen gas and a base:
      • Ca (s) + 2 H 2 O (l)  Ca(OH) 2 (aq) + H 2 (g)
    • 2 valence electrons
    • Form +2 ion by losing those valence electrons
    • Form oxides like CaO, MgO, BaO
  • 67. Alkaline Earth
    • The alkaline earth elements are metallic elements found in the second group of the periodic table. All alkaline earth elements have an oxidation number of +2, making them very reactive.
    • Because of their reactivity, the alkaline metals are not found free in nature.
  • 68. Groups 3-12: Transition Metals
    • Many can form different possible charges of ions
    • If there is more than one ion listed, give the charge as a Roman numeral after the name
    • Cu +1 = copper (I) Cu +2 = copper (II)
    • Compounds containing these metals can be colored.
  • 69. Transition Metals
    • The 38 elements in groups 3 through 12 of the periodic table are called "transition metals". As with all metals, the transition elements are both ductile and malleable, and conduct electricity and heat.
    • The interesting thing about transition metals is that their valence electrons, or the electrons they use to combine with other elements, are present in more than one shell. This is the reason why they often exhibit several common oxidation states.
    • There are three noteworthy elements in the transition metals family. These elements are iron, cobalt, and nickel, and they are the only elements known to produce a magnetic field.
  • 70. Group 17: Halogens
    • Most reactive nonmetals
    • React violently with metal atoms to form halide compounds: 2 Na + Cl 2  2 NaCl
    • Only found in compounds in nature
    • Have 7 valence electrons
    • Gain 1 valence electron from a metal to form -1 ions
    • Share 1 valence electron with another nonmetal atom to form one covalent bond.
  • 71. Halogens
    • The halogens are five non-metallic elements found in group 7 of the periodic table. The term "halogen" means "salt-former" and compounds containing halogens are called "salts". All halogens have 7 electrons in their outer shells, giving them an oxidation number of -1. The halogens exist, at room temperature, in all three states of matter:
    • Solid - Iodine, Astatine
    • Liquid - Bromine
    • Gas - Fluorine, Chlorine
  • 72. Group 18: Noble Gases
    • Are completely nonreactive since they have eight valence electrons, making a stable octet.
    • Kr and Xe can be forced, in the laboratory, to give up some valence electrons to react with fluorine.
    • Since noble gases do not naturally bond to any other elements, one atom of noble gas is considered to be a molecule of noble gas. This is called a monatomic molecule. Ne represents an atom of Ne and a molecule of Ne.
  • 73. Noble Gases
    • The six noble gases are found in group 18 of the periodic table. These elements were considered to be inert gases until the 1960's, because their oxidation number of 0 prevents the noble gases from forming compounds readily.
    • All noble gases have the maximum number of electrons possible in their outer shell (2 for Helium, 8 for all others), making them stable.
  • 74. Diatomic Molecules
    • Br, I, N, Cl, H, O and F are so reactive that they exist in a more chemically stable state when they covalently bond with another atom of their own element to make two-atom, or diatomic molecules.
    • Br 2 , I 2 , N 2 , Cl 2 , H 2 , O 2 and F 2
    • The decomposition of water: 2 H 2 O  2 H 2 + O 2
  • 75. Electronegativity
    • An atom’s attraction to electrons in a chemical bond.
    • F has the highest, at 4.0
    • Fr has the lowest, at 0.7
    • If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion.
    • If the two atoms have an END of less than 1.7, they will share their unpaired valence electrons…covalent bond!
  • 76. Ionization Energy
    • The energy required to remove the most loosely held valence electron from an atom in the gas phase.
    • High electronegativity means high ionization energy because if an atom is more attracted to electrons, it will take more energy to remove those electrons.
    • Metals have low ionization energy. They lose electrons easily to form (+) charged ions.
    • Nonmetals have high ionization energy but high electronegativity. They gain electrons easily to form (-) charged ions when reacted with metals, or share unpaired valence electrons with other nonmetal atoms.
  • 77. Ions
    • Ions are charged particles formed by the gain or loss of electrons.
      • Metals lose electrons (oxidation) to form (+) charged cations .
      • Nonmetals gain electrons (reduction) to form (-) charged anions .
    • Atoms will gain or lose electrons in such a way that they end up with 8 valence electrons (stable octet).
      • The exceptions to this are H, Li, Be and B, which are not large enough to support 8 valence electrons. They must be satisfied with 2 (Li, Be, B) or 0 (H).
  • 78. New Chemical Element
    • Recent hurricanes and gasoline issues are proof of the existence of a new chemical element. Research has led to the discovery of the heaviest element yet known to science. The new element, Governmentium (Gv), has one neutron, 25 assistant neutrons, 88 deputy neutrons, and 198 assistant deputy neutrons, giving it an atomic mass of 312. 
    • These 312 particles are held together by forces called morons, which are surrounded by vast quantities of lepton-like particles called peons. Since Governmentium has no electrons, it is inert; however, it can be detected, because it impedes every reaction with which it comes into contact.  
    • A minute amount of Governmentium can cause a reaction that would normally take less than a second to take from four days to four years to complete. Governmentium has a normal half-life of 2- 6 years; It does not decay, but instead undergoes a reorganization in which a portion of the assistant neutrons and deputy neutrons exchange places. In fact, Governmentium's Mass will actually increase over time, since each reorganization will cause more morons to become neutrons, forming isodopes. This characteristic of moron promotion leads some scientists to believe that Governmentium is formed whenever morons reach a critical concentration. This hypothetical quantity is referred to as critical morass.
    • When catalyzed with money, Governmentium becomes Administratium, an element that radiates just as much energy as Governmentium since it has half as many peons but twice as many morons.
    • Isn't Science Wonderful!!!!!!
  • 79. Now You Should Be Able To…
    • Analyze stable and unstable isotopes of an element
    • to determine the relationship between the isotope’s
    • stability and it application.
    • Describe similarities and differences of isotopes of
    • elements
    • Give examples of uses of isotopes in everyday life
    • • Carbon-12 and its isotopes
    • • Unstable isotopes are radioactive
    • • Isotopes of hydrogen
    • • Define stable and unstable
  • 80. Now You Should Be Able To…
    • Describe the physical and chemical characteristics of
    • an element using the periodic table and make
    • inferences about its chemical behavior.
      • Characteristics:
      • • Metal, nonmetal, metalloid
      • • Reactivity
      • • Valence electrons
      • • Solid, liquid, gas
      • • Oxidation number
      • • Electronegativity/ Electron affinity
    • Periodic Table:
    • Names: Inner transitions, noble gases, transition
    • metals, halogens, alkali metals, alkaline
    • earth metals
      • • Groups and periods
    • Predict properties of elements based on the
    • element’s position on the periodic table