PRINCIPLES OF CHEMISTRY I

CHEM 1211

CHAPTER 8

DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences

Clayton state university

CHAPTER 8

PERIODIC PROPERTIES OF

THE ELEMENTS

- Elements in a given group have similar chemical propertiesbecause the outer-shell electron arrangements are similar

Group IIA elementsBe: 1s22s2

Mg: 1s22s22p63s2

Ca: 1s22s22p63s23p64s2

Sr: 1s22s22p63s23p64s23d104p65s2

CLASSIFICATION OF THE ELEMENTS

- The last electron in an element’s electron configuration causes the difference in the electron configuration of the

preceding element and is referred to as the distinguishing electron

HomeworkWrite notes (one page) on the different classifications of the elements based on electronic properties. Briefly describe the

s-area, p-area, d-area, and the f-area.

CLASSIFICATION OF THE ELEMENTS

CLASSIFICATION OF THE ELEMENTS

Elements can be classified as Metals or Nonmetals

- Based on physical properties

Elements can also be classified as Noble-gas, Representative,

Transition, or Inner Transition- Based on electron configuration

CLASSIFICATION OF THE ELEMENTS

Noble-gas Elements

- Group VIIIA (18) elements on the periodic table (far right column)

- Gases at room temperature

- Little tendency to form chemical compounds

- Electron configuration ends in p6

- Completes p subshell (except Helium)

- Nonmetals

CLASSIFICATION OF THE ELEMENTS

Representative Elements

- Elements in the

s-area (Groups IA and IIA)

first five columns of the p-area (Groups IIIA, IVA, VA, VIA, and VIIA)

- Metals and nonmetals

CLASSIFICATION OF THE ELEMENTS

Transition Elements

- Elements in the d-area of the periodic table

- Groups IIIB (3), IVB (4), VB (5), VIB (6), VIIB (7), VIIIB (8, 9, 10), IB (11), and IIB (12)

- Distinguishing electron in a d subshell

- Metals

CLASSIFICATION OF THE ELEMENTS

Inner Transition Elements

- Elements in the f-area of the periodic table

- The two-row block of elements below the main table

- Distinguishing electron in an f subshell

- Metals

VALENCE ELECTRONS

- Electrons in the highest principal quantum number of an atom, and any electrons in an unfilled subshell

from a lower shell

- Electrons in filled subshells that have lower principal quantum numbers

CORE ELECTRONS

VALENCE ELECTRONS

- Representative elements in the same group of the periodic table have the same number of valence electrons

- The number of valence electrons for representative elements is the same as the group number (with A) in the periodic table

- The maximum number of valence electrons for any given element is eight

VALENCE ELECTRONS

- Not all electrons in a given atom participate in bonding

- Only valence electrons are available for bonding

- For representative and noble-gas elements these electrons are always found in the s or p subshells

VALENCE ELECTRONS

- The electron configuration can be used to determine the number of valence electrons of an atom

C: 1s22s22p2

O: 1s22s22p4

Na: 1s22s22p63s1

As: 1s22s22p63s23p64s23d104p3

ELECTRON CONFIGURATION OF IONS

Cl: [Ne]3s23p5

Cl-: [Ne]3s23p6

Na: [Ne]3s1

Na+: [Ne]

F : 1s22s22p5

F- : 1s22s22p6

Co: [Ar]3d74s2

Co3+: [Ar]3d6

EFFECTIVE NECLEAR CHARGE

-Negatively charged electrons are attracted to the positively charged nucleus

The force of attraction between an electron and the nucleus- Depends on the magnitude of the net nuclear charge acting

on the electrons

- Depends on the average distance between the nucleus and the electron

(Coulombs law)

The force of attraction- Increases with increasing nuclear charge

- Decreases with increasing average distance between electrons and the nucleus

- Electrons also experience repulsion by other electrons in the atom

Zeff = Z – S

Zeff = effective nuclear charge Z = actual nuclear charge (number of protons in the nucleus)(> Zeff)

S = screening constant (represents number of core electrons)

EFFECTIVE NECLEAR CHARGE

Atomic number of Na = 11Number of valence electrons = 1Number of core electrons = 10

As simplified from this modelZ = 11S = 10

Zeff = 11-10 = +1

In actual factZeff in Na is about +2.5

EFFECTIVE NECLEAR CHARGE

In general Zeff in s orbital > Zeff in p orbital > Zeff in d orbital > Zeff in f orbital

- This is the result of the trend in energy levels ns < np < nd < df

- Zeff increases across the periods (from left to right) in the periodic table (Z increases but S remains the same)

- Zeff changes slightly down the groups in the periodic table(essentially the same in a given group)

EFFECTIVE NECLEAR CHARGE

SIZES OF ATOMS

- Atomic radius tends to decrease across the periods (from left to right) in the periodic table

- Due to increase in effective nuclear charge which draws valence electrons closer to the nucleus

- Atomic radius tends to increase down the groups (from top to bottom) of the periodic table

- Due to increase in principal quantum number of the outer electrons (number of shells)

- Cations are smaller than their parent atoms

- Decrease in the number of electrons decreases electron-electron repulsions

- Anions are larger than their parent atoms

- Increase in the number of electrons increases electron-electron repulsions

- Ionic size increases down the group of the periodic table for ions carrying the same charge

SIZES OF ATOMS

ISOELECTRONIC SERIES

- A group of atoms and ions containing the same number of electrons

Due to the same number of electrons- Ionic radius decreases with increasing nuclear charge (electrons are more strongly attracted to the nucleus)

increasing nuclear charge O2- , F- , Na+ , Mg2+ , Al3+

S2- , Cl- , K+ , Ca2+ , Ga3+

Se2- , Br- , Rb+ , Sr2+ , In3+

decreasing ionic radius

IONIZATION ENERGY

- The energy required to remove an electron from a gaseous atom or ion

X(g) → X+(g) + e-

E is positive

- The atom or ion is assumed to be in its ground state

- The highest energy electron is always removed first

Units: kJ/mol96.485 kJ/mol = 1 eV

Ionization Energies of Magnesium (Mg)

Mg(g) → Mg+(g) + e- I1 = 738 kJ/mol

Mg+(g) → Mg2+(g) + e- I2 = 1450 kJ/mol

Mg2+(g) → Mg3+(g) + e- I3 = 7734 kJ/mol

I1 < I2 < I3

IONIZATION ENERGY

First Ionization Energy (I1)- Energy required to remove the highest energy electron of an atom

- The first electron is removed from a neutral atom

- The second electron is removed from a positive ion (more difficult)

- Increase in positive charge binds electrons more tightly

- Large jump is observed in going from removal of valence electrons to removal of core electrons

IONIZATION ENERGY

- First ionization energy increases across the period of the periodic table (from left to right)

- Electrons added in the same principal quantum number do not completely shield increasing nuclear charge

- First ionization energy decreases down the group of the periodic table (from top to down)

- As n increases, the size of orbital increases (distance from nucleus increases) and electrons are easier to remove

IONIZATION ENERGY

- Discontinuities are due to electron repulsions and shielding (Be to B, N to O)

- Representative elements show a larger range of values of I1 than the transition-metal elements

- Smaller atoms have higher ionization energies

IONIZATION ENERGY

- Several cations with the pseudo-noble-gas electron configurationare more stable

[noble gas](n-1)d10

In: [Kr]5s24d105p1

In+: [Kr]5s24d10

Pseudo = In3+: [Kr]4d10

Sn: [Kr]5s24d105p2

Sn2+: [Kr]5s24d10

Pseudo = Sn4+: [Kr]4d10

PSEUDO-NOBLE-GAS CONFIGURATION

ELECTRON AFFINITY

- The energy change that occurs when an electron is added to a gaseous atom

- A measure of the attraction of the atom for the added electron

- Energy is released when an electron is added to most atoms

X(g) + e- → X-(g)

E is negative

Units: kJ/mol

Cl(g) + e- → Cl-(g) E = -349 kJ/mol

- The greater the attraction between an atom and an added electron the more negative the atom’s electron affinity

- Electron affinities for noble gases are positive values (E > 0)

- Halogens have the most negative electron affinities

- Electron affinity changes slightly down the group of the periodic table

ELECTRON AFFINITY

METALS

- Refer to chapter 2 for properties of metals

- Metals tend to have low ionization energies

- Metals form positive ions relatively easily

- Metals lose electrons (oxidize) when they undergo chemical reactions

Charge on metals- Alkali metals: 1+

- Alkaline earth metals: 2+

METALS

- The charge on transition metals do not follow any obvious pattern

- Transition metals are able to form more than one positive ion

- Compounds of metals and nonmetals are ionic

2Na(s) + Cl2(g) → 2NaCl(s) (contains Na+ and Cl- ions)

Most Metal Oxides are Basic

- Dissolve in water to form metal hydroxides

Na2O(s) + H2O(l) → 2NaOH(aq)

O2-(aq) + H2O(l) → 2OH-(aq)(net ionic equation)

- React with acid to form salt and water

NiO(s) + 2HNO3(aq) → Ni(NO3)2(aq) + H2O(l)

METALS

NONMETALS

- Refer to chapter 2 for properties of nonmetals

- Nonmetals tend to have high electron affinities

- Nonmetals tend to gain electrons when they react with metals

- Compounds composed of only nonmetals are generally molecular substances

Most Nonmetal Oxides are Acidic

CO2(g) + H2O(l) → H2CO3(aq)(acidity of rainwater)

- Dissolve in basic solutions to form salt and water

CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l)

NONMETALS

ALKALI METALS (GROUP IA)

- Alkali means ‘ashes”

Relatively abundant in the - earth’s crust (Na, K)

- sea water - human bodies

- Have low densities and melting points

- Very reactive and readily lose an electron to form 1+ ions

- Form hydrides with hydrogen and sulfides with sulfur

2M(s) + H2(g) → 2MH(s) 2M(s) + S(s) → M2S(s)

- React vigorously with water to produce hydrogen gas and alkali metal hydroxide

(very exothermic and may explode)

2M(s) + 2H2O(l) → 2MOH(aq) + H2(g)

ALKALI METALS (GROUP 1A)

- Can react with oxygen to form oxides, peroxides, and superoxides

Li forms only oxides (O2-)4Li(s) + O2(g) → 2Li2O(s) (lithium oxide)

Na can form peroxides (O2

2-)2Na(s) + O2(g) → Na2O2(s) (sodium peroxide)

K, Rb, and Cs can form peroxides (O22-) and superoxides (O2

-)K(s) + O2(g) → KO2(s) (potassium superoxide)

ALKALI METALS (GROUP 1A)

ALKALINE EARTH METALS (GROUP 2A)

Compared to alkali metals- harder

- more dense - higher melting points

- less reactive than the respective adjacent alkali metal

- Tend to lose two outer s electrons to from 2+ ions

- Give off characteristic colors when heated in a flame (salts used in fireworks)

- Reactivity increases down the group

- Beryllium does not react with water

- Magnesium reacts slowly with water

- Calcium and elements below react readily with water Ca(s) + 2H2O(l) → Ca(OH)2 + H2(g)

ALKALINE EARTH METALS (GROUP 2A)

HYDROGEN

- First element in the periodic table (1s1 electron configuration)

- Nonmetal

- Can be metallic under extreme pressures

- Colorless diatomic gas

- Has very high ionizaton energy - More than double those of alkali metals

- Due to absence of nuclear shielding of the 1s electron

- Does not easily lose its valence electron

- Share with nonmetals to form molecular compounds

- Can lose its electron to form a cation (H+)

- Can gain electron to form the hydride ion (H-)

HYDROGEN

CHALCOGENS (GROUP 6A)THE OXYGEN GROUP

- Properties change from nonmetallic to metallic down the group Nonmetallic properties: oxygen, sulfur, selenium

Metallic properties: tellurium and below

- Oxygen is a colorless gas at room temperature

- The other group members are solids

- Oxygen exists as O2 (oxygen gas) and O3 (ozone) - Allotropes (different forms of the same element in the same state)

- O2 can produce O3 in lightning storms

3O2(g) → 2O3(g) Ho = 284.6 kJ

- Sulfur also has several allotropic forms - The most common is S8 (yellow solid)

CHALCOGENS (GROUP 6A)THE OXYGEN GROUP

THE HALOGENS (GROUP 7A)

- All halogens are nonmetals

- Melting and boiling points increase with increasing atomic number

- Consist of diatomic molecules (F2, Cl2, Br2, and I2)

- Form colored gases

THE HALOGENS (GROUP 7A)

At room temperature - Fluorine and chlorine are gases

- Bromine is a liquid - Iodine is a solid

- Have highly negative electron affinities

- Tend to gain electrons to form 1- ions

- Reactivity decreases down the group

- React readily with most metals to form ionic halides

- React with hydrogen to form gaseous hydrogen halides H2(g) + X2 → 2HX(g)

- Hydrogen halides dissolve in water to form acids [HCl(aq)]

- Fluorine is very reactive (dangerous)

- Chlorine is the most industrially useful

THE HALOGENS (GROUP 7A)

THE NOBLE GASES (GROUP 8A)

- Nonmetals

- Monatomic

- Gases at room temperature

- Have completely filled s and p subshells

- Have high first ionization energies

- Have stable electron configuration

- Unreactive