ii. atom. periodic table and trends ii. atom. periodic table and trends prepared by phd halina...

II. Atom. Periodic Table and Trends

II. Atom. Periodic Table and Trends

Prepared by PhD Halina Falfushynska

Rutherford model (planetary model) of atom

Draw a diagram showing the location of each part of the atom.

The components that make up the atom are known as  sub-atomic particles. Some particles have charges which are integer multiples of the elementary charge. 

Components of Atoms 

Quark

• Quarks combine to form composite particles called hadrons, the most stable of which are protons and neutrons.

• Quarks have various intrinsic properties, including electric charge, color charge, mass, and spin.

For comparison: 

• Human hair  ~106  atoms wide •   Pencil line  ~106  atoms wide •   HIV virus    ~800 atoms wide (~108  atoms total) •    E. coli    ~1011  atoms total 

Bohr Model of AtomBohr Model of Atom• Electrons in atoms orbit the nucleus.• The electrons can only orbit stably, without 

radiating, in certain orbits (called by Bohr the "stationary orbits“. These orbits are associated with definite energies and are also called energy shells.

Electrons can only gain and lose energy by jumping from one allowed orbit to another, absorbing or emitting electromagnetic radiation

Bohr Model• Since the energy states are quantized, the light emitted 

from excited atoms must be quantized and appear as line spectra.

• After lots of math, Bohr showed that

where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else).

Line Spectra and the Bohr ModelLine Spectra and the Bohr Model

2

18 1J 10178.2n

En

Bohr Model

• We can show that

• nf is the final energy level, and ni is the initial energy level

• When ni > nf, energy is emitted.

• When nf > ni, energy is absorbed

Line Spectra and the Bohr ModelLine Spectra and the Bohr Model

22

18 11J 10178.2

if

fi nn

hchE

hEEE if

2

18 1J 10178.2n

En

Rydberg formula

Bohr Model

Line Spectra and the Line Spectra and the Bohr ModelBohr Model

• Schrödinger proposed an equation that contains both wave and particle terms.

• Solving the equation leads to wave functions. • The wave function gives the shape of the electronic 

orbital.  [“Shape” really refers to density of electronic charges.]

• The square of the wave function,  gives the probability of finding the electron ( electron density ).

Quantum Mechanics and Atomic OrbitalsQuantum Mechanics and Atomic Orbitals

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Quantum Mechanics and Atomic OrbitalsQuantum Mechanics and Atomic Orbitals

Solving Schrodinger’s Equation gives rise to ‘Orbitals.’  

These orbitals provide the electron density distributed about the nucleus.

Orbitals are described by quantum numbers.

Orbitals and Quantum Numbers• Schrödinger’s equation requires 3 quantum numbers:

1. Principal Quantum Number, n.  This is the same as Bohr’s n.  As n becomes larger, the atom becomes larger and the electron is further from the nucleus. ( n = 1 , 2 , 3 , 4 , …. )

2. Angular Momentum Quantum Number, .  This quantum number depends on the value of n.  The values of   begin at 0 and increase to (n - 1).  We usually use letters for   (s, p, d and f for   = 0, 1, 2, and 3).  Usually we refer to the s, p, d and f-orbitals. 

3. Magnetic Quantum Number, m.  This quantum number depends on  .  The magnetic quantum number has integral values between -   and +  .  Magnetic quantum numbers give the 3D orientation of each orbital.

Quantum Mechanics and Atomic OrbitalsQuantum Mechanics and Atomic Orbitals

Quantum Numbers of WavefuntionsQuantum Numbers of Wavefuntions

Quantum # Symbol Values Description

Principal n 1,2,3,4,… Size & Energy of orbital

Angular Momentum

0,1,2,…(n-1)

for each n

Shape of orbital

Magnetic m -…,0,…+ for each

Relative orientation of orbitals within same

Spin ms +1/2 or –1/2 Spin up or Spin down

Angular Momentum Quantum # () Name of Orbital

0 s (sharp)

1 p (principal)

2 d (diffuse)

3 f (fundamental)

4 g

The s-Orbitals

Representations of OrbitalsRepresentations of Orbitals

The p-Orbitals

Representations of OrbitalsRepresentations of Orbitals

d-orbitals

Many-Electron Atoms Many-Electron Atoms

Orbitals and Their Energies

Electron Spin and the Pauli Exclusion Principle

Many-Electron Atoms Many-Electron Atoms

Electron Spin and the Pauli Exclusion Principle

• Since electron spin is quantized, we define ms = spin quantum number =  ½.

• Pauli’s Exclusions Principle:: no two electrons can have the same set of 4 quantum numbers.• Therefore, two electrons in the same orbital must have 

opposite spins.

Many-Electron Atoms Many-Electron Atoms

Hund's rules• For a given electron configuration, the term with 

maximum multiplicity has the lowest energy. The multiplicity is equal to, where  is the total spin angular momentum for all electrons. 

• For a given multiplicity, the term with the largest value of the total orbital angular momentum quantum number  has the lowest energy.

• For a given term, in an atom with outermost subshell half-filled or less, the level with the lowest value of the total angular momentum quantum number  (for the operator ) lies lowest in energy. 

Figure 6.27

Orbitals CD

Atomic and Mass Numbers, Isotopes

• Using the periodic table, the atomic number (number of protons; used to identify the element), elemental symbol, and atomic weight can be identified. 

Nuclides• Atoms may exist as more than one isotope – 

atoms with the same number of protons but have a different number of neutrons. The mass number is used to distinguish isotopes.

• One of two or more atoms whose nuclei have the same number of neutrons but different numbers of protons  are called isotones.

• One of two or more atoms or elements having the same atomic weights or mass numbers but different atomic numbers are called isobars

Average atomic weight• The average atomic weight is a weighted

average of naturally occurring isotopes and fractional abundance in nature.

Example: Carbon has two appreciably present isotopes, 12C and 13C.

Respectively, the abundances are 98.9% and 1.1%. • average atomic weight = (0.989)(12 amu) +

(0.011)(13 amu) = 12.011 amu (on periodic table)

av. atomic wt. = (mass isotope A)(% A) + (mass isotope B)(% B)

Molecules and Molecular Compounds

• Sharing electrons molecular (covalent) compound

• Trading electrons ionic compound

ChargesCharges

• A cation is formed when more protons than electrons are present in an atom or molecule

• An anion is formed when more electrons than protons are present in an atom or molecule.

Examples

F199 Na23

11 O188

+ 2-

9 protons 10 neutrons 9 electrons

11 protons 12 neutrons 10 electrons

8 protons 10 neutrons 10 electrons

Predicting Ionic Charges

• Metals form cations. • Nonmetals form anions.

Naming Inorganic Compounds• Positive ions: • •  Cations formed from metal (main group or transition) atoms have the same name. 

• •  If a metal can form different cations, the positive charge is indicated by a Roman numeral in parenthesis following the name of the metal. Older names using  –ous and –ic are still seen but their use is fading. 

• •  Cations formed from nonmetals have names that end in –ium (i.e. hydronium ion and ammonium ion).

The Periodic Law says:• When elements are arranged in order

of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

• Horizontal rows = periods–There are 7 periods

• Vertical column = group (or family)–Similar physical & chemical prop.–Identified by number & letter (IA, IIA)

Areas of the periodic table• Three classes of elements are: 

1) Metals: electrical conductors, have luster, ductile, malleable

2) Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity

3) Metalloids: border the line-2 sides– Properties are intermediate between 

metals and nonmetals

Electron Configurations in Groups

• Elements can be sorted into 4 different groupings based on their electron configurations:

1) Noble gases2) Representative elements3) Transition metals4) Inner transition metals

Classify elements based on electron configuration.

• Group IA – alkali metals–Forms a “base” (or alkali) when

reacting with water (not just dissolved!)

• Group 2A – alkaline earth metals–Also form bases with water; do not

dissolve well, hence “earth metals”• Group 7A – halogens

–Means “salt-forming”

Electron Configurations in Groups

1) Noble gases are the elements in Group 8A  (also called Group18 or 0)

– Previously called “inert gases” because they rarely take part in a reaction; very stable = don’t react

– Noble gases have an electron configuration that has the outer s and p sublevels completely full

Electron Configurations in Groups

2) Representative Elements are in Groups 1A through 7A

– Display wide range of properties, thus a good “representative” 

– Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquids

– Their outer s and p electron configurations are NOT filled

Electron Configurations in Groups

3) Transition metals are in the “B” columns of the periodic table

– Electron configuration has the outer s sublevel full, and is now filling the “d” sublevel

– A “transition” between the metal area and the nonmetal area

– Examples are gold, copper, silver

Electron Configurations in Groups

4) Inner Transition Metals are located below the main body of the table, in two horizontal rows

– Electron configuration has the outer s sublevel full, and is now filling the “f” sublevel

– Formerly called “rare-earth” elements, but this is not true because some are very abundant

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10 5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s24f14

5d106p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

Do you notice any similarity in these configurations of the alkali metals?

ALL Periodic Table Trends • Influenced by three factors:

1. Energy Level–Higher energy levels are further away 

from the nucleus.2. Charge on nucleus (# protons)–More charge pulls electrons in closer. 

(+ and – attract each other)• 3. Shielding effect

(blocking effect?)

What do they influence?

Energy levels and Shielding have 

an effect on the GROUP ( )

Nuclear charge has an effect on a 

PERIOD ( )

#1. Atomic Size - Group trends

• As we increase the atomic number (or go down a group).

• each atom has another energy level,

• so the atoms get bigger.

HLi

Na

K

Rb

#1. Atomic Size - Period Trends

• Going from left to right across a period, the size gets smaller.

• Electrons are in the same energy level.• But, there is more nuclear charge.• Outermost electrons are pulled closer.

Na Mg Al Si P S Cl Ar

Ions

• Some compounds are composed of particles called “ions”–An ion is an atom (or group of atoms) 

that has a positive or negative charge• Atoms are neutral because the number of 

protons equals electrons–Positive and negative ions are formed 

when electrons are transferred (lost or gained) between atoms

Ions

• Metals tend to LOSE electrons, from their outer energy level, and thus a positively charged particle is formed = “cation”. Cations are smaller than the atom they came from 

• Nonmetals tend to GAIN one or more electrons. Negative ions are called 

“anions”. Anions are bigger than the atom they came from 

Ion Group trends• Each step down a 

group is adding an energy level

• Ions therefore get bigger as you go down, because of the additional energy level.

Li1+

Na1+

K1+

Rb1+

Cs1+

Ion Period Trends• Across the period from left to right, the 

nuclear charge increases - so they get smaller.• Notice the energy level changes between 

anions and cations. (more protons would pull the same # of electrons in closer)

Li1+

Be2+

B3+

C4+

N3- O2- F1-

#2. Trends in Ionization Energy

• Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).

• Removing one electron makes a 1+ ion.• The energy required to remove only 

the first electron is called the first ionization energy.

Ionization Energy• The second ionization energy is the 

energy required to remove the second electron.–Always greater than first IE.

• The third IE is the energy required to remove a third electron.–Greater than 1st or 2nd IE.

Ionization Energy - Group trends

• As you go down a group, the first IE decreases because...–The electron is further away 

from the attraction of the nucleus, and

–There is more shielding.

Ionization Energy - Period trends

• All the atoms in the same period have the same energy level.

• Same shielding.• But, increasing nuclear charge• So IE generally increases from left to 

right.• Exceptions at full and 1/2 full orbitals.

Electron AffinitiesElectron Affinities• Electron affinity is the opposite of ionization 

energy.• Electron affinity: the energy change when a 

gaseous atom gains an electron to form a gaseous ion:

Cl(g) + e-  Cl-(g)• Electron affinity can either be exothermic (as the 

above example) or endothermic:Ar(g) + e-  Ar-(g)