The Structure of the AtomThe AtomAtoms are basic building blocks ofmatter, and cannot be chemicallysubdivided by ordinary means.
in 1808 John Dalton (an Englishscientist) states a theory about thenature of the elements known asDalton’s atomic theory , the mainideas of this theory can be statedas follows:
Elements are made of a tiny particlecalled an atom. All atoms of a given element areidentical.The atoms of a given element aredifferent from those of any otherelement.
Atoms of one element can combinewith atoms of other elements to form compounds. In chemical reactions atoms are neither created nor destroyed they simply change the way they are grouped together.
Later on Thomson & Rutherformed worked on the structure of the atom & they discovered that the atom isconsist of a tiny nucleus ( about 10-13 cm indiameter) and electrons that move around thenucleus . The nucleus contains protons whichhave a positive charge equal in magnitude tothe electrons negative charge , and neutrons ,which have almost the same mass as a protonbut no charge.
Atoms are composed of three type of particles: Protons Neutrons Electron
The mass & charge of the electron, proton & neutron are given below: particle Relative Relative mass charge electron 1 1- proton 1836 1+ neutron 1839 No charge
Modern atomic theory: In 1911 Niels Bohr construct a model of the hydrogen atom with quantized energy levels , Bohr picture the electrons moving in circular orbits (like planetorbiting the sun) , corresponding to thevarious allowed energy levels. He suggestedthat the electron could jump to different orbitby absorbing or emitting energy.
Bohr’s atomic orbital:• Is a specific path on which the electrons travel about the nucleus. - - - +
Although Bohrs model opened the way for the later theories it isimportant to realize that electrons donot move around the nucleus incircular orbits like the planet orbitingthe sun.
Later on Schrodinger found that it is not precisely to describe theelectrons path , he could only predictthe probability of finding the electronat a given point in space around thenucleus .
• In its ground state the hydrogenelectron has a probability map The probability map, or orbital that describes the hydrogen electron in its lowest possible energy state . The more intense the color of agiven dot the more likely it is that theelectron will be found at that point.
Shrodinger showed that the orbitals of electrons are regions of electron density with the location and routsof electrons described asprobabilities.
Isotopes:Atoms with the same number ofprotons but different number ofneutrons. In nature elements areusually found as a mixture ofisotopes.
Hydrogen 1 (hydrogen) Hydrogen 2 (deuterium) Hydrogen 3 (tritium) 1 proton, 0 neutrons 1 proton, 1 neutron 2 neutrons Mass number = 1 Mass number = 2 Mass number = 3
ExampleThree isotopes of elemental carbonare C612 , C613, C614 . Determine thenumber of each of the three types subatomic particles in each of thesecarbon atoms.
Atomic number:The number of protons in the nucleus.Mass number:The sum of the number of neutrons &number of protons in a given nucleus.
Ions: • Under certain circumstances it is possible to remove electrons from a neutral atom leaving a positively charged particle (cation) . • Electrons may be added to a certain atoms to form a negatively charged particle (anion). These charged particles whether positive or negative are called ions.
Atomic OrbitalQuantum numbers: The various orbitals available to an atom are described by four quantum numbers, which can take certain values to create differently sized and shaped orbital of various energies:
The principal quantum number (n)describes the: Shells:(an electron shell is collection of orbitals) Size of orbital Numbered 1, 2, 3, 4, 5, etc are often lettered (K, L, M, etc.).
The subsidiary quantum number (l)describes the: Subshells : are groups of orbitals with in an electron Shape of orbital (number of lobes). Given letters s, p, d, f, g, h, i, etc. The values of ( l ) run from 0 to n − 1
The magnetic quantum number (m)describes: Orientation of the orbitals in space. Named after the directions they point in (x, y, z, etc.) Can also be given numbers ranging from 0, ±1, ±2, ±3 … ±l.
The spin quantum number (s). describes: spin of an electron on its own axis May have the values +½ or -½. Two electrons in the same orbital have paired (opposite) spins
The Quantum Numbersname symbol valuesPrincipal QuantumNumber n any integer from 1 to infinitySubsidiary QuantumNumber l any integer from 0 to n-1Magnetic QuantumNumber m any integer from - l to+ lSpin QuantumNumber s 1/2 -/+
orbital Max no. ofL-value No. of orbitals type electrons 0 s 1 2 1 p 3 6 2 d 5 10 3 f 7 14
Using symbols, the valid quantumstates can be listed in the followingmanner: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 6p 6d 6f 6g 7h 7s 7p 7d 7f 7g 7h 8i
Atomic Orbitals The S Orbital The simplest orbital in the atom is the 1s orbital. The 1s orbital is simply a sphere of electron density. There is only one s orbital per shell The s orbital can hold two electrons have different spin quantum numbers l = 0 m = 0 S = +1/2 , -1/2
The P Orbitals Starting from the 2nd shell, there is a set of p orbitals There are 3 choices for the magnetic quantum number, which indicates 3 differently orientated p orbitals px, py, and pz each orbital can accommodate two electrons, giving a total capacity of 6 electrons. p orbitals are very often involved in bonding
Electronic configuration of elementsRules for electronic configuration: The Pauli Exclusion principle: Only two electrons with opposite spin can occupy an atomic orbital (or no twoelectrons have the same (4) quantumnumbers n, l, m and s) 1s
Hunds Rule: Electrons prefer parallel spins in separate orbitals of subshells The Aufbau Principle: Explains the order in which the electrons fill the various orbitals in an atom. Filling begins with the orbitals in the lowest- energy shells and continues through thehigher-energy shells
The energy relationships among the first threelevels of orbitals;
Special electronic configuration: The pairing of electrons raise the orbital energy slightly. Half-filled and full-filled subshell low the energy. For example the electronic configuration of Cr and 1sCu 2s 2p 3s 3p 4s 3dCr24 1s 2s 2p 3s 3p 4s 3dCu29
Tutorial (A)which of the following pair of atoms contain the same -1 ?number of neutronsA- C614 & C6 12B- F919 & Ne1022C- S1632 &Al1329
Which of the following particles does not contain -2)- the same number of electrons as fluoride ion )FA- Ne B- + Li+ C- Na
the electronic configuration of helium atom ,boron -3atom , carbon atom & element X are given below.Which one could be the electronic configuration ofelement X?A- 1s2 2s22P1B- 1s2C- 1s2 2s 1D- 1s2 2s2 2p2
4- atoms which have the same electronicconfiguration are said to be isoelectronic whichof the following is not isoelectronic with O2-? A- N 3- B- Al 3+ C- Na + D- Na
the magnetic quantum number )m( gives -5A- the sub shellsB- the orbitals .C- the spin of the electron
Tutorial (B)1- an orbital can take a maximum of…………..electrons.2- an) s ( sub level can take a maximum of……electrons.3- the )p( sub level can take a maximum of…….electrons.4- how many electrons can fit in to a set of )5d( orbitals?5- what is the maximum numbers of electrons in levels, 2, 3, 4 ?
which of the following orbitals could not be exist? -6. 1s,1p,1d,2s,2p,2d,3p,3d,3f,4s,4f 7- write the four quantum numbers for energy level)4(? 8- For the H-like atom, which subshell has the highest energy level? 4f, 3d, 2p, 1s
9- How many electrons are required to fill all the following subshells? 1s 2s 2p 3s 3p 4s 3d 4p10- Which of the following two electronic configuration is more stable? a- [Ar]4s1 3d5 b- [Ar]4s2 3d4
11- Which of the following two electronicconfigurations is more stable? a [Ar]4s2 3d9 b [Ar]4s1 3d1012- Choose the electronic configuration for palladium, Pd )Z = 46(. a- [Kr]5s1 4d7 b- [Kr]5s1 4d8 c- [Kr]5s0 4d10 d- [Kr]5s1 4d10
The Periodic Table of Elements By the late 1800s many elements had already been discovered. The scientist Dmitri Mendeleev, a Russian chemist, proposed an arrangement of known elements based on their atomic mass The modern arrangement of the elements is known as the Periodic Table of Elements and is arranged according to the atomic number of elements.
Periodic law:The modern periodic law states that ; when elementsare arranged by atomic number; their physical andchemical properties vary periodically .
The periodic table displays the elements in rows )periods ( and columns in order of increasingatomic number. Elements that have similar chemical properties fall into vertical columns called groups or families Most of the elements are metals and located on the left hand side of the periodic table The nonmetals appear on the right hand side of the periodic table
From the periodic table we can know many informationdirectly like symbol, atomic number, atomic mass & youcan also know whether the element is metal, non metalor metalloid.Todays periodic table consist of seven horizontal rowscalled periods & a number of vertical columns calledgroups.
All elements in each group have the same number ofelectrons in their outer most shells so the behave similarly.We have two types of groups:Group )A(: ) representative or main elements(Group )B(:) transition elements(
Metals : are found on the left-hand and at the centre ofthe periodic tableNon metals: are relatively few they are in the upper-righthand corner the tableMetalloids : those are few elements exhibit both metallicand nonmetallic behavior )also known as semimetals(
Group 1A metals )Li, Na, K, Rb, Cs, and Fr( are calledthe alkali metals , and they are the most reactive metalsin the periodic table.Group 2A metals )Be, Mg, Ca, and Ra( are called thealkaline earth metals.
Group 1B )Cu, Ag and Au( are called the coinage metalsAlthough these elements have outer electronicconfiguration similar to those of the alkali metals and thealkaline earth metal the are much less reactiveGroup 7A )F, Cl, Br and I( are called halogens and theyare the most reactive nonmetals in the periodic tableGroup 8A )He, Ne, Ar, Kr, Xe, and Rn( are called therare or noble gases,these gases are characteristic by their completely filledshells, for this reason they are chemically inert.
Physical properties of elementsIonization Energy )I(:Is the energy required to remove an electron from anisolated atom in its ground state.X)g(
Chemical Bonds:In nature elements or compounds exist due tocombination of similar or different atomsAtoms and molecules are electrically neutral, accordingto this fact in atom combination each tries to exhibit 8electrons on the outer most shell that by loosing,gaining or sharing electrons.
Metals: have three or less electrons e.g. Na, Ca, Fe Nonmetals: Have five or more electrons e.g. H, N, S Carbon atom: have four electrons in the outer most shell liable to loose them or to gain more electrons
The Bonds By which atoms can combine are : Ionic Bond Covalent Bond Co-ordinate Bond
Ionic BondThis is Characteristic of metallic and non metalliccombination forming a neutral molecule. It is achieved via two steps: Ionization of atoms Formation of the molecule
Ionization of atoms: + -Na Na + e Cations ++ -Ca Ca + 2e - - Cl + e Cl Anions -- S + 2e S
Formation of the molecule In formation of molecules, loss or gain of electron occurs at the same time.
e.g. Formation of NaCl: + Na - Na + e Chlorine atom attacks this electron to its outermost shell and become an anion. - - Cl + e Cl
Thus two ions of different charges of equal numberareformed, attraction between them taken place, leadingto the formation of the neutral molecule NaCl. + - Na + Cl NaCl NaCl is said to be ionic compound
Example:Formation of CaCl2 ; ++ - Ca Ca + 2e - - Cl + e Cl - ++ Cl Ca + - CaCl2 Cl
Covalent Bond The covalent bond is the chemical bond in which two or more non metal atoms share electrons Both atoms are unable to loose or gain electrons By sharing electrons both atoms reach octet state
e.g. Fluorine atom has 7 electrons in its outermost shell. It needs one electron to reach its octet. To achieve this F atom shares an electron belonging to another F atom By this F2 molecule is formed