Atomic Structure.

Atomic Structure

Atomic Structure All matter is composed of atoms.
Understanding the structure of atoms is critical to understanding the properties of matter. properties of solid materials depend on the geometrical atomic arrangements, and the interactions between constituent atoms.

(greek for indivisible)
history of the atom 460 BC Democritus developed the idea of atoms he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA (greek for indivisible) It took ~2400 years from when it was conceived to the time experimental evidence prove of the atom existence.

history of the atom ATOMS 1808 John Dalton
suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

history of the atom ELECTRON 1898 Joseph John Thompson / Cambridge
found that atoms could sometimes eject a far smaller negative particle which he called ELECTRON 1906 Nobel prize in Physics

History of the atom 1910 Ernest Rutherford / Cambridge student of Thompson proposed a more detailed model with a central nucleus: positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction 1908 Nobel prize in Chemistry

History of the atom 1913 Niels Bohr / Danish / a football fanatic
studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons. 1922 Nobel prize in physics

Bohr’s atom Rutherford’s model predicted a rainbow of colors rather than discrete lines obtained from an atomic line spectra. To explain the line spectra, Bohr proposed that electrons of specific energy moved in circular orbits around the nucleus and could not exist between these orbits.

Atomic Structure Atoms are composed of
protons – positively charged particles neutrons – neutral particles electrons – negatively charged particles in orbitals surrounding the nucleus. nucleus

Atomic Structure Every different atom has a characteristic number of protons in the nucleus. atomic number (Z) = number of protons For an electrically neutral atom, atomic number = number of electrons. Atoms with the same atomic number have the same chemical properties and belong to the same element.

Atomic Structure Z ranges from 1 for hydrogen to 92 for uranium (the highest for the naturally occuring elements).

Atomic Structure The atomic mass (A) of a specific atom: the sum of the number of protons and neutrons within the nucleus. mass number: A = Z + N number of protons is the same for all atoms of a given element, number of neutrons (N) may be variable.

Atomic Structure The number of protons in the nucleus of the atom is equal to the atomic number (Z). The number of electrons in a neutral atom is equal to the number of protons. The mass number of the atom (M) is equal to the sum of the number of protons and neutrons in the nucleus. The number of neutrons is equal to the difference between the mass number of the atom (M) and the atomic number (Z).

Atomic Structure / isotopes
atoms of some elements have two or more different atomic masses, called isotopes.

Atomic Structure Atomic weight: Weighted average of the atomic masses of the atom’s naturally occurring isotopes. Boron consists of the isotopes: 19.7% B-10 and 80.3% B-11. atomic mass A, B = (19.7 x 10)+(80.3 x 11)]/100= 10.8 Bromine’ isotopes: 50.5% Br-79 and 49.5% Br-81. atomic mass A, Br=[(50.5 x 79)+(49.5 x 81)]/100= 80.0

Atomic Structure In one mole of a substance there are x (Avogadro’s number) atoms or molecules. Atomic weight = weight of x 1023 atoms For example, the atomic weight of iron is amu/atom, or g/mol. 1 amu/atom = 1g/mol = 1 dalton Atomic mass unit (amu): 1⁄12 of the atomic mass of the carbon C: / H:1.008 etc.

subatomic particles particle Mass (g) Charge (C/eV) Electron (e-)
9.11x10-28 -1.6x10-19 -1 Proton (p) 1.67x10-24 +1.6x10-19 +1 Neutron (n) Proton is 1837 times heavier than an electron. Neutron is 1842 times heavier than an electron. Electron is much lighter with respect to the protons and neutrons

Atomic Structure 4 He 2 HELIUM ATOM + - + - proton Shell nucleus
# electrons = # protons electron neutron He 2 4 ATOMIC MASS NUMBER = number of protons + number of neutrons ATOMIC NUMBER = number of protons

Lithium 7 Li 3 Electrons Protons Neutrons three electrons
three protons four neutrons. 3

Beryllium 9 Be 4 Electrons Protons Neutrons four electrons
four protons five neutrons. 4

Boron 11 B 5 Electrons Protons Neutrons five electrons five protons
six neutrons. 5

Carbon 12 C 6 Electrons Protons Neutrons six electrons six protons
six neutrons. 6

Nitrogen 14 N 7 Electrons Protons Neutrons seven electrons
seven protons seven neutrons. 7

Oxygen 16 O 8 Electrons Protons Neutrons eight electrons eight protons
eight neutrons 8

Fluorine 19 F 9 Electrons Protons Neutrons nine electrons nine protons
ten neutrons. 9

Neon 20 Ne 10 Electrons Protons Neutrons ten electrons ten protons
ten neutrons 20 Ne 10

Sodium 23 Na 11 Electrons Protons Neutrons eleven electrons
eleven protons twelve neutrons 23 Na 11

Atomic structure 1 56 27 H Fe Al 1 26 13 48 39 238 Ti K U 22 19 92
How many protons, neutrons and electrons? 1 56 27 H Fe Al 1 26 13 48 39 238 Ti K U 22 19 92

Atomic structure How many protons, neutrons and electrons? Ti 22 48

Atomic structure How many protons, neutrons and electrons? 26 56 Fe

Atomic structure How many protons, neutrons and electrons? Al 13 27

Atomic structure How many protons, neutrons and electrons? K 19 39

Atomic structure The charge and mass number of an electron are:
charge = 0, Mass number = 1 charge = -1, Mass number = 0 charge = +1, Mass number = 1 charge = +1, Mass number = 0 The charge and mass number of a neutron are?

Atomic structure Which of the following has 25 protons and 31
neutrons? 56Mn 56Ga 25Ga 31Mn 56Ba

Atomic structure Why does chlorine have an atomic mass of 35.5, which is not a whole number? Chlorine contains an extra electron which makes it weigh more than 35. Chlorine contains 17 protons and 18.5 neutrons Chlorine normally exists in an excited state, and so it weighs more than 35. The chlorine was not pure when its atomic mass was measured. Chlorine, as found in nature, contains a mixture of the isotopes 35Cl and 37Cl, in such proportions as to give an average atomic mass of 35.5

Atomic structure The two main parts of an atom are?
a) nucleus and electron energy levels b) nucleons and protons c) oxidation number and valence d) protons and neutrons e) protons and electrons

Atomic structure The nucleus of the element having atomic number 25 and atomic weight 55 will contain? 25 protons and 30 neutrons 30 protons and 25 neutrons 55 protons 55 neutrons

Atomic structure A beryllium atom has 4 protons, 5 neutrons, and 4 electrons. What is the mass number of this atom? 4 5 8 9 13

Atomic structure The smallest particle into which an element can be divided and still have the properties of that element a) nucleus b) electron c) atom d) neutron How would you describe the nucleus? a) dense, positively charged b) mostly empty space, positively charged c) tiny, negatively charged d) dense, negatively charged

Atomic structure Where are electrons likely to be found?
a) in the nucleus b) in electron clouds c) mixed throughout an atom d) in definite paths Every atom of a given element has the same number of a) protons b) neutrons c) electrons d) isotopes

Atomic structure What is the meaning of the word atom? a) dividable
b) invisible c) hard particles d) not able to be divided Which statement is true about isotopes of the same element? a) They have the same number of protons b) They have the same number of neutrons c) They have a different atomic number d) They have the same mass

Atomic structure Which has the least mass in an atom? a) nucleus
b) proton c) neutron d) electron If an isotope of uranium, uranium-235, has 92 protons, how many protons does the isotope uranium-238 have? a) 92 b) 95 c) 143 d) 146

Atomic structure A carbon atom with 6 protons, 6 electrons, and 6 neutrons would have a mass number of a) 6 b) 12 c) 15 d) 18 The number at the top is the a) atomic number b) element name c) atomic mass d) chemical symbol

Atomic structure How many electrons does a neutral Cl atom contain? 16 17 18 19 What is the difference between atomic mass and atomic weight? Cl 17 35

Atomic Structure Neutral atoms have the same number
of protons and electrons. Ions are charged atoms. cations – have more protons than electrons and are positively charged anions – have more electrons than protons and are negatively charged

Atomic Structure Na Na+ Cl- Cl e- + + e-
If a neutral atom looses one or more electrons it becomes a cation. Na 11 protons 11 electrons Na+ 11 protons 10 electrons e- + If a neutral atom gains one or more electrons it becomes an anion. Cl- 17 protons 18 electrons Cl 17 protons 17 electrons + e-

Bohr Atomic model electrons are assumed to revolve around the atomic nucleus in discrete orbitals, and the position of any particular electron is more or less well defined in terms of its orbital. Electrons are permitted to have only specific values of energy.

Bohr Atomic model

excitation vs relaxation
An electron may change energy by making a quantum jump either to an higher energy (with absorption of energy) or to a lower energy (with emission of energy). relaxation excitation

Quantum Mechanics Unfortunately, extremely small particles (electrons) do not follow the laws of classical (Newtonian) physics. The new physics that mathematically treats small particles is called Quantum Mechanics.

electron distribution
wave-mechanical model an electron is no longer treated as a particle moving in a discrete orbital; electron is considered to exhibit both wavelike and particle-like characteristics. The position of an electron is described by a probability distribution // electron cloud.

Quantum Mechanics Wave behavior is described with the wave function ψ, incorporating the wave and particle features of electrons (Erwin Schrödinger) The probability of finding an electron in a certain area of space is proportional to ψ2 electron density. Austrian; 1933 Nobel prize in physics

Quantum Mechanics

Heisenberg’s uncertainty principle
more precisely the position of some particle is determined, the less precisely its momentum can be known A macroscale analogy… High Shutter Speed Low Shutter Speed Can judge location, Can judge speed, but not speed. But not location

Heisenberg’s uncertainty principle
we cannot precisely measure the momentum and the position of an electron at the same time. As the momentum of the electron is more and more certain, the position of the electron becomes less and less certain, and vice versa. n = 2.5 cannot exist as a principal quantum number. There must be an integral number of wavelengths (n) in order for an electron to maintain a standing wave. If there were to be partial waves, the whole and partial waves would cancel each other out and the particle would not move.

Quantum Mechanics The Schrödinger equation specifies possible energy states an electron can occupy. The energy states and wave functions are characterized by a set of quantum numbers. Instead of orbits in the Bohr model, quantum numbers and wave functions describe atomic orbitals.

quantum numbers every electron in an atom is characterized by four quantum numbers. There are three quantum numbers necessary to describe an atomic orbital. The principal quantum number (n) designates size The angular moment quantum number (l) describes shape The magnetic quantum number (ml) specifies orientation

Principal Quantum Number (n)
designates the size of the orbital. Larger values of n correspond to larger orbitals. The allowed values of n are integers: 1, 2, 3 and so forth. A collection of orbitals with the same value of n is frequently called a shell. n K L M N O P 1 2 3 4 5 6 ……

Angular moment Quantum Number (l)
signifies the subshell describes the shape of the orbital. l values range from 0 to n – 1 Example: If n = 2, l can be 0 or 1. n 1 2 3 4 5 6 l subshell s p d f g h energy state 7 9 11

Magnetic Quantum Number (ml)
describes the orientation of the orbital in space. ml are integers that depend on l: – l,…0,…+l ml identifies # of energy states for each subshell For an s subshell: a single energy state For p, d, and f subshells: 3, 5, and 7 energy states

Number of available electron states for initial shells and subshells
Principal Quantum No: n Shell Subshell l No. of energy States: ml Number of Electrons Per Subshell Per Shell 1 K s /0 1 / 0 2 L s / 0 8 p / 1 3 / -1,0,+1 6 3 M 18 d / 2 5 / -2,-1,0,+1,+2 10 4 N 32 f / 3 7 / -3,-2,-1,0,+1,+2,+3 14 5 O 50 7 / -3,-2,-1,0,1,2,3 g / 4 9 / -4,-3,-2,-1,0,+1,+2,+3,+4 P 72 9 / -4,-3,-2,-1,0,1,2,3,4 h / 5 11 / -5,-4,-3,-2,-1,0,+1,+2,+3,+4,+5 22

Atomic orbitals An s subshell has one orbital which is spherically shaped. If you were to measure where the electron was within an s subshell many, many times and plot the results on a graph you would get something like this.

l = 1 (as required for a p orbital)
Atomic Orbitals p orbitals: a dumbbell shape with electrons on either side of the nucleus in tear drop shaped lobes Three orientations: l = 1 (as required for a p orbital) ml = –1, 0, +1

l = 2 (as required for a d orbital)
Atomic Orbitals The d orbitals: Five orientations: l = 2 (as required for a d orbital) ml = –2, –1, 0, +1, +2

l = 3 (as required for a d orbital)
Atomic orbitals d-orbitals are followed by the seven f-orbitals. 7 orientations: l = 3 (as required for a d orbital) ml = -3, –2, –1, 0, +1, +2, +3

2px Quantum Numbers To summarize quantum numbers: principal (n) – size
angular (l) – shape magnetic (ml) – orientation electron spin (ms) direction of spin Required to describe an atomic orbital principal (n = 2) 2px related to the magnetic quantum number (ml ) angular momentum (l = 1) Required to describe an electron in an atomic orbital

Electron Spin Quantum Number (ms)
used to specify an electron’s spin. There are two possible directions of spin. Allowed values of ms are +½ and −½.

Pauli exclusion principle
No Two Electrons in an Atom Can Have the Same Four Quantum Numbers; the same values for n, l, ml, and ms. Although the first three quantum numbers identify a specific orbital and may have the same values, the fourth is significant and must have opposite spins. a set of quantum numbers is specific to a certain electron.

Quantum numbers An electron with n = 2, ℓ = 1, ml = −1, and ms = +1/2
is found in the same atom as a second electron with n = 2, ℓ = 1, and ml = −1. What is the spin quantum number for the second electron?

Quantum numbers First electron: n = 1, ℓ = 1, ml = −1, ms = +1/2
Second electron: n = 1, ℓ = 1, ml = −1, ms = ? Since the first three quantum numbers are identical for these two electrons, we know that they are in the same orbital. As a result, the spin quantum number for the second electron cannot be the same as the spin quantum number for the first electron. This means that the spin quantum number for the second electron must be ms = −1/2.

Quantum numbers An electron with n = 5, ℓ = 4, ml = 3, and ms = −1/2 is found in the same atom as a second electron with n = 5, ℓ = 4, and ml = 3. What is the spin quantum number for the second electron?

Quantum numbers First electron: n = 5, ℓ = 4, ml = 3, ms = −1/2
Second electron: n = 5, ℓ = 4, ml = 3, ms = ? Since the first three quantum numbers are identical for these two electrons, we know that they are in the same orbital. As a result, the spin quantum number for the second electron cannot be the same as the spin quantum number for the first electron. This means that the spin quantum number for the second electron must be ms = +1/2.

Quantum numbers Can an electron with
n = 1, ℓ = 0, ml = 0, and ms = +1/2 exist in the same atom as a second electron with n = 2, ℓ = 0, ml = 0, and ms = +1/2?

Pauli exclusion principle
First electron: n = 1, ℓ = 0, ml = 0, ms = +1/2 Second electron: n = 2, ℓ = 0, ml = 0, ms = +1/2 Since these two electrons are in different orbitals, they occupy different regions of space within the atom. As a result, their spin quantum numbers can be the same, and thus these two electrons can exist in the same atom.

Atomic structure Maximum number of electrons in a subshell with l = 3 and n = 4 is 10 12 14 16 18

Atomic structure The lowest principal quantum number for an electron is? 1 2 3 4

Atomic structure Which sublevel can by occupied by a maximum of 10 electrons? s p d f

Atomic structure The K, L and M shells of an atom are full. Its atomic number is_______. 18 20 10 12

Atomic structure State whether an electron can be described by each of the following sets of quantum number. If a set is not possible, state why not. n = 2, l = 1, ml = -1 n = 1, l = 1, ml = +1 n = 4, l = 3, ml = +3 n = 3, l = 1, ml = -3

Atomic structure State whether an electron can be described by each of the following sets of quantum number. If a set is not possible, state why not. n = 2, l = 1, ml = -1 l: 0/1, ml: -1/0/1 n = 1, l = 1, ml = +1 l: 0, ml: 0 n = 4, l = 3, ml = +3 l: 0,1,2,3; ml: -3,-2,-1,0,1,2,3 n = 3, l = 1, ml = -3 l:0,1,2; ml: -2,-1,0,1,2

Atomic structure n = 3, l = 1, ml = ? n = 4, l = ?, ml = -2
Replace the question marks by suitable responses in the following quantum number assignments. n = 3, l = 1, ml = ? n = 4, l = ?, ml = -2 n = ?, l = 3, ml = ?

Atomic structure n = 3, l = 1, ml = -1,0,1 n = 4, l = 2, ml = -2
Replace the question marks by suitable responses in the following quantum number assignments. n = 3, l = 1, ml = -1,0,1 n = 4, l = 2, ml = -2 n = 4, l = 3, ml = -3,-2,-1,0,1,2,3

Atomic structure Provide the three quantum numbers describing each of the three p orbitals in the 2p subshell. n l ml 2px 2py 2pz

Atomic structure n = 1 l = 0 ml= 0 1 ml= -1,0,1
For n = 1, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml. n = 1 l = 0 ml= 0 1 ml= -1,0,1

Atomic structure How many orbitals in shell n = 1?
How many electrons possible?

Atomic structure How many orbitals in shell n = 2? How many electrons possible? For n = 2, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml.

Atomic structure 2 L 2 L 8 Principal Quantum No: n Shell Subshell l
No. of energy States: ml 2 L s / 0 1 / 0 p / 1 3 / -1,0,+1 Principal Quantum No: n Shell Subshell l Number of Electrons Per Subshell Per Shell 2 L s / 0 8 p / 1 6

Atomic structure For n = 3, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml. How many orbitals in shell n = 2? How many electrons possible?

Atomic structure 3 M 3 M 18 Principal Quantum No: n Shell Subshell l
No. of energy States: ml 3 M s / 0 1 / 0 p / 1 3 / -1,0,+1 d / 2 5 / -2,-1,0,+1,+2 Principal Quantum No: n Shell Subshell l Number of Electrons Per Subshell Per Shell 3 M s / 0 2 18 p / 1 6 d / 2 10

Atomic structure For n = 4, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml.

Atomic structure 4 N Principal Quantum No: n Shell Subshell l
No. of energy States: ml 4 N s / 0 1 / 0 p / 1 3 / -1,0,+1 d / 2 5 / -2,-1,0,+1,+2 f / 3 7 / -3,-2,-1,0,+1,+2,+3

Atomic structure n l ml ms
Provide the four quantum numbers describing each of the two electrons in the 3s orbital. n l ml ms

Atomic structure n l ml ms 3 0 0 +1/2 3 0 0 -1/2
Provide the four quantum numbers describing each of the two electrons in the 3s orbital. n l ml ms /2 /2

Quantum Numbers: A Macroscale Analogy
n - indicates which train (shell) l - indicates which car (subshell) ml - indicates which row (orbital) ms - indicates which seat (spin) No two people can have exactly the same ticket (sit in the same seat).

Electron energy states
electrons have discrete energy states they fill up the lowest possible energy states in the electron shells and subshells, When all the electrons occupy the lowest possible energies in accord with the foregoing restrictions, an atom is said to be in its ground state. Energy states for a Na atom

Electron configurations
Most elements: Electron configuration not stable! Electron configuration (stable) ... 1s 2 2s 2p 6 3s 3p 3d 10 4s 4p Atomic # 18 36 Element 1 Hydrogen Helium 3 Lithium 4 Beryllium 5 Boron Carbon Neon 11 Sodium 12 Magnesium 13 Aluminum Argon Krypton

Electron Configurations
Valence electrons – those in unfilled shells Filled shells more stable Valence electrons are most available for bonding and tend to control the chemical, electrical, thermal, optical properties example: C (atomic number = 6) 1s2 2s2 2p2 valence electrons

the full electronic configuration of an element is 1s22s22p5 How many electrons does it have in its out shell? the full electronic configuration of an element. What is its atomic number?

Electronic Configurations
Fe-atomic # = 26 1s2 2s2 2p6 3s2 3p6 3d 6 4s2 1s 2s 2p K-shell n = 1 L-shell n = 2 3s 3p M-shell n = 3 3d 4s 4p 4d Energy N-shell n = 4

Valence electrons Element symbol Atomic number e-configuration
# of valence electrons H 1 1s1 He 2 1s2 Li 3 1s22s1 Be 4 1s22s2 B 5 1s22s22p1 C 6 1s22s22p2 N 7 1s22s22p3 O 8 1s22s22p4 F 9 1s22s22p5 Ne 10 1s22s22p6

Order of Subshell Filling
The electron configurations of the first ten elements illustrate this point.

Electron configurations for common elements

Shells and subshells In multi-electron atoms, the energies of the atomic o orbitals are split. Splitting of energy levels refers to the splitting of a shell (n=3) into subshells of different energies (3s, 3p, 3d)

Splitting of Shells into subshells
3s subshell (n = 3; l = 0) 3rd shell (n = 3) 3p subshell (n = 3; l = 1) 3d subshell (n = 3; l = 2) 2s subshell (n = 2; l = 0) 2nd shell (n = 2) 2p subshell (n = 2; l = 1)

rules for electron configurations: Electrons will reside in the lowest possible energy orbitals Each orbital can accommodate a maximum of two electrons. Electrons will not pair in degenerate orbitals if an empty orbital is available. Orbitals will fill in the order ..3p6/4s2/3d10/4p6/5s2/4d10/ 5p6/6s2/4f14/5d10/6p6/7s2

Energy Level Diagram of a Many-Electron Atom
6s p d f 32 5s p d 18 4s p d Arbitrary Energy Scale 18 3s p 8 Original reference: Pimental, Chemistry An Experimental Science, (CHEM Study), 1969, page 266. 2s p 8 1s 2 NUCLEUS

The electron configuration describes how the electrons are distributed in the various atomic orbitals. In a ground state hydrogen atom, the electron is found in the 1s orbital. Ground state electron configuration of hydrogen principal (n = 1) 1s1 number of electrons in the orbital or subshell Energy 2s 2p 2p 2p angular momentum (l = 0) The use of an up arrow indicates an electron with ms = + ½ 1s

If hydrogen’s electron is found in a higher energy orbital, the atom is in an excited state. Energy 2s 2p 2p 2p A possible excited state electron configuration of hydrogen 1s 2s1

Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers. The ground state electron configuration of helium Energy 2p 2p 2p 1s2 2s Quantum number Principal (n) Angular moment (l) Magnetic (ml) Electron spin (ms) 1 1 1s describes the 1s orbital describes the electrons in the 1s orbital + ½ ‒ ½

The Aufbau principle states that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals. Li has a total of 3 electrons The ground state electron configuration of Li 1s22s1 Energy 2p 2p 2p 2s The third electron must go in the next available orbital with the lowest possible energy. 1s The 1s orbital can only accommodate 2 electrons (Pauli exclusion principle)

Be has a total of 4 electrons Energy 2p 2p 2p 2s The ground state electron configuration of Be 1s 1s22s2

B has a total of 5 electrons Energy 2p 2p 2p 2s The ground state electron configuration of B 1s 1s22s22p1

Hund’s rule, the most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized. C has a total of 6 electrons The ground state electron configuration of C 1s22s22p2 Energy 2p 2p 2p 2s The 2p orbitals are of equal energy, or degenerate. Put 1 electron in each before pairing (Hund’s rule). 1s

N has a total of 7 electrons Energy 2p 2p 2p 2s The 2p orbitals are of equal energy, or degenerate. Put 1 electron in each before pairing (Hund’s rule). 1s 1s22s22p3 The ground state electron configuration of N

O has a total of 8 electrons Energy 2p 2p 2p 2s Once all the 2p orbitals are singly occupied, additional electrons will have to pair with those already in the orbitals. 1s The ground state electron configuration of O 1s22s22p4

F has a total of 9 electrons Energy 2p 2p 2p 2s When there are one or more unpaired electrons, as in the case of oxygen and fluorine, the atom is called paramagnetic. 1s The ground state electron configuration of F 1s22s22p5

Ne has a total of 10 electrons Energy 2p 2p 2p 2s When all of the electrons in an atom are paired, as in neon, it is called diamagnetic. 1s 1s22s22p6 The ground state electron configuration of Ne

Worked Example Write the electron configuration and give the orbital diagram of a calcium (Ca) atom (Z = 20). Z = 20, Ca has 20 electrons. Each s subshell can contain a maximum of two electrons, whereas each p subshell can contain a maximum of six electrons. Solution Ca 1s22s22p63s23p64s2 1s2 2s2 2p6 3s2 3p6 4s2 Remember that the 4s orbital fills before the 3d orbitals.

Worked Example electron configuration for an arsenic atom (Z = 33) in the ground state. Z = 18 for Ar. The order of filling beyond the noble gas core is 4s, 3d, and 4p. Fifteen electrons go into these subshells because there are 33 – 18 = 15 electrons in As beyond its noble gas core. 2 2 6 2 6 10 Solution As [Ar]4s23d104p3 2 3 Arsenic is a p-block element; therefore, we should expect its outermost electrons to reside in a p subshell.

electron configuration?
Al 13 27 Number of Energy Levels: 3 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 3 1s22s22p63s23p1

K 19 39 Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 8 Fourth Energy Level: 1 1s22s22p63s23p64s1

22 48 Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 10 Fourth Energy Level: 2 1s22s22p63s23p64s23d2

Cr 24 52 Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 13 Fourth Energy Level: 1 1s22s22p63s23p64s13d5

Mn 25 55 Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 13 Fourth Energy Level: 2 1s22s22p63s23p64s23d5

26 56 Fe Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 14 Fourth Energy Level: 2 1s22s22p63s23p64s23d6

Cu 29 64 Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 18 Fourth Energy Level: 1 1s22s22p63s23p64s13d10

Zn 30 64 Number of Energy Levels: 4 First Energy Level: 2 Second Energy Level: 8 Third Energy Level: 18 Fourth Energy Level: 2 1s22s22p63s23p64s23d10

Valence electrons They occupy the outermost shell.
They participate in the bonding between atoms They dictate the physical and chemical properties if the outermost or valence electron shell are completely filled: stable electron configurations occupation of the s and p states for the outermost shell by a total of eight electrons, in neon (Ne), argon (Ar), and krypton (Kr); inert, or noble, gases, which are virtually unreactive chemically.

Valence electrons unfilled valence shells assume stable electron configurations by gaining or losing electrons to form charged ions, or by sharing electrons with other atoms. This is the basis for some chemical reactions, and also for atomic bonding in solids

Valence electrons Under special circumstances, the s and p orbitals combine to form hybrid spn orbitals, where n indicates the number of p orbitals involved, which may have a value of 1, 2, or 3. The IIIA, IVA, and VA group elements of the periodic table often form these hybrids. The driving force for the formation of hybrid orbitals is a lower energy state for the valence electrons. For carbon the sp3 hybrid is of primary importance in organic and polymer chemistries.

‘Aufbau’ Principle: filling orbitals
s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p n = 1 l = 0 ml = 0 n = 2 ml = 1 ml = -1 H: 1s1

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p n = 1 l = 0 ml = 0 n = 2 ml = 1 ml = -1 He: 1s2

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p n = 1 l = 0 ml = 0 n = 2 ml = 1 ml = -1 Li: 1s2 2s1

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p n = 1 l = 0 ml = 0 n = 2 ml = 1 ml = -1 Be: 1s2 2s2

s- and p-orbitals B: 1s2 2s22p1 ‘Aufbau’ Principle: filling orbitals
‘core’ closed shell open shell: valence electrons

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p C: 1s2 2s22p2 Hund’s rule: maximum number of unpaired electrons is the lowest energy arrangement.

Hund’s rule electrons fill orbitals one at a time.
we must fill each shell with one electron each before starting to pair them up. the charge of an electron is negative and electrons repel each other. An electron will try to create distance between itself and other electrons by staying unpaired. This further explains why the spins of electrons in an orbital are opposite (i.e. +1/2 and -1/2).

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p N: 1s2 2s22p3 O: 1s2 2s22p4

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals 1s 2s 2p F: 1s2 2s22p5 Ne: 1s2 2s22p6

s- and p-orbitals ‘Aufbau’ Principle: filling orbitals Na: 1s22s22p63s1 or [Ne]3s1 P: [Ne]3s23p3 Ar: [Ne]3s23p6 Mg: 1s22s22p63s2 or [Ne]3s2

Which one of the following is a proper orbital configuration?

beyond the d-orbitals ‘s’-groups ‘p’-groups group
lanthanides actinides ‘s’-groups ‘p’-groups d-transition elements f-transition elements group period 1s2 2s2/2p6 3s2/3p6/ 4s2/3d10/4p6 5s2/4d10/5p6/ 6s2/4f14/5d10/6p6

Organisation of the periodic table

Give electron configurations for the Fe3+and S2- ions.

The Fe3+ ion is an iron atom that has lost three electrons. Since the electron configuration of the Fe atom is 1s22s22p63s23p63d64s2, the configuration for Fe3+ is 1s22s22p63s23p63d5. The S2- ion a sulfur atom that has gained two electrons. Since the electron configuration of the S atom is 1s22s22p63s23p4, the configuration for S2- is 1s22s22p63s23p6.

Give the electron configurations for the following ions? Fe2+ [Ar] 3d64s electrons Al3+ [Ne] 3s23p electrons Cu+ [Ar] 3d104s1 -1 electron Ba2+ [Xe] 6s electrons Br- [Ar] 3d104s24p5 +1 electron O2- [He] 2s22p electrons

Which of the following electron configurations is an inert gas, a halogen, an alkali metal, an alkaline earth metal, a transition metal? a) 1s22s22p63s23p63d74s2 b) 1s22s22p63s23p6 c) 1s22s22p5 d) 1s22s22p63s2 e) 1s22s22p63s23p63d24s2 f) 1s22s22p63s23p64s1

The halogens (group 7A, or group 17, of the periodic table) all have similar chemical properties (for example, forming singly charged negative ions). What aspect of their electron configurations leads to these elements having such similarities? They all have a complete 1s2 shell at the lowest energy level They all have an identical Ns2Np5 configuration for their valence electrons (N is any whole number). They all have p electrons in their outermost shell. They all have an odd number of protons. They all have an even number of neutrons.

Atomic Structure.

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Atomic Structure.

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