1 2 Periodic Trends in Atomic Properties 3 Characteristic properties and trends of the elements are...

2

Periodic Trends in Periodic Trends in Atomic PropertiesAtomic Properties

3

Characteristic properties and trends of the elements are the basis of the periodic table’s design.

4

These trends allow us to use the periodic table to accurately predict properties and reactions of a wide variety of substances.

5

Metals and NonmetalsMetals and NonmetalsMetals and NonmetalsMetals and Nonmetals

6

Chemical Properties of Metals

• metals tend to lose electrons and form positive ions (cations).

• nonmetals tend to gain electrons and form negative ions (anions).

Chemical Properties of Nonmetals

When metals react with nonmetals electrons are usually transferred from the metal to the nonmetal.

7

Physical Properties of Metals

• lustrous

• malleable

• good conductors of heat

• good conductors of electricity

• nonlustrous

• brittle

• poor conductors of heat

• poor conductors of electricity

Physical Properties of Nonmetals

8

Metalloids have properties that are intermediate between metals and nonmetals

9

The Metalloids

1. boron

2. silicon

3. germanium

4. arsenic

5. antimony

6. tellurium

7. polonium

10

Metals are found to the left of the metalloidsNonmetals are found to the right of the metalloids.

11.1

11

Atomic RadiusAtomic RadiusAtomic RadiusAtomic Radius

12

Radii of atoms increase down a group.

For each step down a group, electrons enter the next higher energy level.

13

Radii of atoms tend to decrease from left to right across a period.

For representative elements within the same period the energy level remains constant as electrons are added.

Each time an electron is added a proton is a added to the nucleus.

This increase in positive nuclear charge pulls all electrons closer to the nucleus.

14

Ionization EnergyIonization EnergyIonization EnergyIonization Energy

15

The ionization energy of an atom is the energy required to remove the outermost electron from an atom.

Na + ionization energy → Na+ + e-

18

Periodic relationship of the first ionization energy for representative elements in the first four periods.

Ionization energies gradually increase from left to right across a period.

IA

IIAIIIA

IVA

VAVIA

VIIA

Noble Gases

1

2

3

4

19

Periodic relationship of the first ionization energy for representative elements in the first four periods.

11.3

Ionization energies of Group A elements decrease from top to bottom in a group.

IA

IIAIIIA

IVA

VAVIA

VIIA

Noble Gases

Distance of O

uter Shell E

lectrons From

Nucleus

nonmetals

metals

nonmetals have higher ionization potentials than metals

20

Lewis “Dot” Structures of AtomsLewis “Dot” Structures of Atoms

21

Metals form cations and nonmetals form anions to attain a stable valence electron structure.

22

This stable structure often consists of two s and six p electrons.These rearrangements occur by losing, gaining, or sharing electrons.

23

• Na with the electron structure 1s22s22p63s1

has 1 valence electron.

The Lewis structure of an atom is a representation that shows the valenceelectrons for that atom.

• Fluorine with the electron structure 1s22s22p5

has 7 valence electrons

25

The Lewis structure of an atom uses dots to show the valence electrons of atoms.

The number of dots equals the number of s and p electrons in the atom’s outermost shell.

BPaired electrons

Unpaired electron

Symbol of the element

2s22p1

26

The number of dots equals the number of s and p electrons in the atom’s outermost shell.

S3s23p4

The Lewis structure of an atom uses dots to show the valence electrons of atoms.

27

Lewis Structures of the first 20 elements.

28

The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

29

With the exception of helium, this structure consists of eight electrons in the outermost energy level.

30

The Covalent Bond:The Covalent Bond: Sharing ElectronsSharing Electrons

31

A covalent bond consists of a pair of electrons shared between two atoms.

In the millions of chemical compounds that exist, the covalent bond is the predominant chemical bond.

32

Substances which covalently bond exist as molecules.

Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms.

33

The term molecule is not used when referring to ionic substances.

Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of NaCl exist.

34

Covalent bonding in the hydrogen molecule

Two 1s orbitals from each of two hydrogen atoms overlap.

Each 1s orbital contains 1 electron.

The orbital of the electrons includes both hydrogen nuclei.

The most likely region to find the two electrons is between the two nuclei.

The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together.

Two 1s orbitals from each of two hydrogen atoms overlap.

36

hydrogen chlorine iodine nitrogen

Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element.

A dash may replace a pair of dots.

H-H

37

ElectronegativityElectronegativity

Linus Pauling

38

electronegativity The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

39

• If the two atoms that constitute a covalent bond are identical then there is equal sharing of electrons.

• This is called nonpolar covalent bonding.

• Ionic bonding and nonpolar covalent bonding represent two extremes.

40

• If the two atoms that constitute a covalent bond are not identical then there is unequal sharing of electrons.

• This is called polar covalent bonding.

• One atom assumes a partial positive charge and the other atom assumes a partial negative charge.– This charge difference is a result of the

unequal attractions the atoms have for their shared electron pair.

41

Polar and Non-Polar

42

:H Cl+ -

Shared electron pair.

:The shared electron pair is closer to chlorine than to hydrogen.

Partial positive charge on hydrogen.

Partial negative charge on chlorine.

Chlorine has a greater attraction for the shared electron pair than hydrogen.

Polar Covalent Bonding in HCl

The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity.

43

A scale of relative electronegativities was developed by Linus Pauling.

44

Electronegativity decreases down a group for representative elements.

Electronegativity generally increases left to right across a period.

45

The electronegativities of the metals are low.The electronegativities of the nonmetals are high.

52

A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points.

A dipole can be written as + -

53

An arrow can be used to indicate a dipole.

The arrow points to the negative end of the dipole.

H Cl H Br H

O

H

Molecules of HCl, HBr and H2O are polar .

54

A molecule containing different kinds of atoms may or may not be polar depending on its shape.

The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

56

Lewis Structures ofLewis Structures ofCompoundsCompounds

57

In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

58

• The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion.

• In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

59

Cl2O has two possible arrangements.

Cl-Cl-OThe two chlorines can be bonded to each other.

Cl-O-ClThe two chlorines can be bonded to oxygen.

Usually the single atom will be the central atom.

60

Practice WritingLewis StructuresPractice WritingLewis Structures

61

Atom GroupValence

Electrons

Cl VIIA 7

H IA 1

C IVA 4

N VA 5

S VIA 6

P VA 5

I VIIA 7

Valence Electrons of Group A Elements

62

3-Dimensional Shapes

Linear180

Bent105

TrigonalPlanar120

Tetrahedral109.5

TrigonalPyramidal

107

63

Covalent Bonding StructuresMolecular Formula

Lewis “dot” Structure

3-D Structure

StructureName

Bond Angle

Polar orNon-polar

H2O

CO2

PH3

NO3–

CH4

Bent 105 Polar

Linear 180 Non-Polar

TrigonalPyramidal

107 Polar

TrigonalPlanar

120 Non-Polar

Tetrahedral 109.5 Non-Polar

64

The Ionic Bond: Transfer ofThe Ionic Bond: Transfer ofElectrons From One AtomElectrons From One Atom

to Anotherto Another

65

After sodium loses its 3s electron it has attained the same electronic structure as neon.

66

After chlorine gains a 3p electron it has attained the same electronic structure as argon.

67

Formation of NaClFormation of NaCl

68

The 3s electron of sodium transfers to the half-filled 3p orbital of chlorine.

Lewis representation of sodium chloride formation.

A sodium ion (Na+) and a chloride ion (Cl-) are formed.

The force holding Na+ and Cl- together is an ionic bond.

69

Formation of MgCl2Formation of MgCl2

70

Two 3s electrons of magnesium transfer to the half-filled 3p orbitals of two chlorine atoms.A magnesium ion (Mg2+) and two chloride ions (Cl-) are formed.The forces holding Mg2+ and two Cl- together are ionic bonds.

71

NaCl is made up of cubic crystals.In the crystal each sodium ion is surrounded by six chloride ions.

72

In the crystal each chloride ion is surrounded by six sodium ions.

73

The ratio of Na+ to Cl- is 1:1

There is no molecule of NaCl

77

• Metals usually have one, two or three electrons in their outer shells.

• When a metal reacts it:– usually loses one two or three electrons

– attains the electron structure of a noble gas

– becomes a positive ion.

• The positive ion formed by the loss of electrons is much smaller than the metal atom.

78

• Nonmetals are usually only a few electrons short of having a noble gas structure.

• When a nonmetal reacts it:– usually gains one two or three electrons

– attains the electron structure of a noble gas

– becomes a negative ion.• The negative ion formed by the gain of

electrons is much larger than the nonmetal atom.

80

Predicting Formulas ofPredicting Formulas ofIonic CompoundsIonic Compounds

81

In almost all stable chemical compounds of representative elements, each atom attains a noble gas electron configuration.

82

Ions are always formed by adding or removing electrons from an atom.

83

Most often ions are formed when metals combine with nonmetals.

•Metals will lose electrons to attain

a noble gas configuration.

•Nonmetals will gain electrons to attain

a noble gas configuration.

84

The charge on an ion can be predicted from its position in the periodic table.

85

elements of Group IIA have a

+2 charge

elements of Group IA have a +1 charge

elements of Group VA have a

-3 charge

elements of Group VIA have a

-2 charge

elements of Group VIIA have a

-1 charge

86

Writing Formulas From Writing Formulas From Names of CompoundsNames of Compounds

87

A chemical compound must have anet charge of zero.

88

If the compound contains ions, then the charges on all of the ions must add to zero.

89

Write the formula of calcium chloride.

Step 1. Write down the formulas of the ions.

Ca2+ Cl-

Step 2. Combine the smallest numbers of Ca2+ and Cl- so that the sum of the charges equals zero.

(2+) + 2(1-) = 0

The correct formula is CaCl2

The lowest common multiple of +2 and –1 is 2

The cation is written first.

The anion is written second.

(Ca2+) + 2(Cl-) = 0

90

Write the formula of barium phosphide.


Ba2+ P3-

Step 2. Combine the smallest numbers of Ba2+ and P3- so that the sum of the charges equals zero.

3(2+) + 2(3-) = 0

The correct formula is Ba3P2


3(Ba2+) + 2(P3-) = 0

The cation is written first.

The anion is written second.

91

Write the formula of magnesium oxide.


Mg2+ O2-

Step 2. Combine the smallest numbers of Mg2+ and O2- so that the sum of the charges equals zero.

(2+) + (2-) = 0

The correct formula is MgO


(Mg2+) + (O2-) = 0

92

Write the Formula of Sodium Peroxide

1 22Na O+ −

gives orNa2O2 NaO

93

Write the Formula of Sodium Peroxide

1 22Na O+ −

NaO does not contain the peroxide anion

gives notNa2O2 NaO

Don’t mess with the subscripts of polyatomic ions!!

101

The Following Slidesare for Extra Practice/

Explanation

The Following Slidesare for Extra Practice/

Explanation

102

Procedures for WritingLewis Structures

Procedures for WritingLewis Structures

103

Step 1 Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion.

–If you are writing the structure of an ion, add one electron for each negative charge or subtract one electron for each positive charge on the ion.

104

Step 1. The total number of valence electrons is eight, two from the two hydrogen atoms and six from the oxygen atom.

Write the Lewis structure for H2O.

105

Step 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (two dots or one dash).– Hydrogen, which contains only one bonding

electron, can form only one covalent bond.

– Oxygen atoms normally have a maximum of two covalent bonds (two single bonds, or one double bond).

106

Step 2. The two hydrogen atoms are connected to the oxygen atom. Write the skeletal structure:


Place two dots between the hydrogen and oxygen atoms to form the covalent bonds.

H O HorH OH

: : : :

107

Step 3. Subtract two electrons for each single bond you used in Step 2 from the total number of electrons calculated in Step 1.

– This gives you the net number of electrons available for completing the structure.

108

Step 3. Subtract the four electrons used in Step 2 from eight to obtain four electrons yet to be used.


H O H: :

109

Step 4. Distribute pairs of electrons (pairs of dots) around each atom (except hydrogen) to give each atom a noble gas configuration.

110

Step 4. Distribute the four remaining electrons in pairs around the oxygen atom. Hydrogen atoms cannot accommodate any more electrons.


These arrangements are Lewis structures because each atom has a noble gas electron structure.

H O HorH OH

: : : ::: ::

The shape of the molecule is not shown by the Lewis structure.

111

Step 1. The total number of valence electrons is 16, four from the C atom and six from each O atom.

Write a Lewis structure for CO2.

112

Step 2. The two O atoms are bonded to a central C atom. Write the skeletal structure and place two electrons between the C and each oxygen.

O C O: :


113


Step 3. Subtract the four electrons used in Step 2 from 16 (the total number of valence electrons) to obtain 12 electrons yet to be used.

O C O: :

114

O C O: :

Step 4. Distribute the 12 electrons (6 pairs) around the carbon and oxygen atoms. Three possibilities exist.

Many of the atoms in these structures do not have eight electrons around them.


O C O: : O C O: :: :::

:: ::: :

:

:4

electrons6

electrons6

electrons6

electrons

::

::: :

6electrons

I II III

115


O C O: :: :::

::

Step 5. Remove one pair of unbonded electrons from each O atom in structure I and place one pair between each O and the C atom forming two double bonds.

O C O::: : ::

O C O::: : : :::::

Each atom now has 8 electrons around it.

Carbon is sharing 4 electron pairs.

double bond double bond

116

Complex Lewis StructuresComplex Lewis Structures

117

There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written.

118

Step 1. The total number of valence electrons is 24, 5 from the nitrogen atom and 6 from each O atom, and 1 from the –1 charge.

Write a Lewis structure for NO2.-3NO .

119

Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge.

Step 2. The three O atoms are bonded to a central N atom. Write the skeletal structure and place two electrons between each pair of atoms.


O N O: :O:

-

120

Step 3. Subtract the 6 electrons used in Step 2 from 24, the total number of valence electrons, to obtain 18 electrons yet to be placed.

Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge.

O N O: :O:

-


121

O N OO

Step 4. Distribute the 18 electrons around the N and O atoms.


:: : :::

:: :: ::

electron deficient

122


Step 5. One pair of electrons is still needed to give all the N and O atoms a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond.

:: : :::

:: :O N O

O: ::

-

123


Step 5. One pair of electrons is still needed to give all the N and O atoms a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond.

N

O::

O:: :

O:::

:-

124

A molecule or ion that shows multiple correct Lewis structures exhibits resonance.


Step 5. There are three possible Lewis structures.

N

O

::O:: :

O:

:

::

-:

N

O

::O:

:

O::

::

-

Each Lewis structure is called a resonance structure.

N

O::

O:: :

O:::

:-

132

Molecular ShapeMolecular Shape

133

The 3-dimensional arrangement of the atoms within a molecule is a significant determinant of molecular interactions.

134

(bent)

135

The Valence ShellThe Valence ShellElectron Pair (VSEPR) Electron Pair (VSEPR)

ModelModel

136

The VSEPR model is based on the idea that electron pairs will repel each other electrically and will seek to minimize this repulsion.

To accomplish this minimization, the electron pairs will be arranged as far apart as possible around a central atom.

137

BeCl2 is a molecule with only two pairs of electrons around beryllium, its central atom. Its electrons are arranged 180o apart for maximum separation.

138

• BF3 is a molecule with three pairs of electrons around boron, its central atom.

• Its electrons are arranged 120o apart for maximum separation.

• This arrangement of atoms is called trigonal planar.

139

• CH4 is a molecule with four pairs of electrons around carbon, its central atom.

• An obvious choice for its atomic arrangement is a 90o angle between its atoms with all of its atoms in a single plane.

• However, since the molecule is 3-dimensional the molecular structure is tetrahedral with a bond angle of 109.5o.

140

Ball and stick models of methane, CH4, and carbon tetrachloride, CCl4.

11.13

141

• Ammonia, NH3, has four electron pairs around nitrogen.

The arrangement of electron pairs around nitrogen is tetrahedral.

142

The NH3 molecule is pyramidal.

NH3 has one unbonded pair of electrons.

143

• Water has four electron pairs around oxygen.

The arrangement of electron pairs around oxygen is tetrahedral.

144

The H2O molecule is bent.

H2O has two unbonded pairsof electrons.

1 2 Periodic Trends in Atomic Properties 3 Characteristic properties and trends of the elements are...

Documents

Transcript of 1 2 Periodic Trends in Atomic Properties 3 Characteristic properties and trends of the elements are...