IB Chemistry · Ions •When an atom gains or loses an electron, it is called an ion. (The valence...

IB Chemistry

Chapter 4.1

Chemical Bonds

• Atoms or ions that are strongly attached to one another

• Chemical bonds will form if potential energy decreases (becomes more stable)

2

Valence Electrons

• Valence electrons are the electrons in the outer shell that are on the same period or row. These are the electrons that are involved in bonding atoms together.

• Valence electrons can easily be found by finding the group number the element is in

Nitrogen has 5 valence electrons (it is in group 5) Group 1 has 1, group 2 has 2, group 3 has 3, etc.

Ions

• When an atom gains or loses an electron, it is called an ion. (The valence electrons are the ones lost)

• Cation if positively charged or giving electrons away.

Anion if negatively charged or accepting electrons.

Ions

• Metals tend to lose their valence electrons when bonding and become cations.

This is due to lower ionization energies

• Nonmetals tend to gain electrons in their highest energy level when bonding and become anions. This is due to high electronegativity and high ionization energies

• The most stable atoms have 8 electrons in their highest energy level (the octet rule)

Predicting Ionic Charges

Group 1: Lose 1 electron to form 1+ ions

H+ Li+ Na+ K+


Group 2: Loses 2 electrons to form 2+ ions

Be2+ Mg2+ Ca2+ Sr2+ Ba2+

Predicting Ionic Charges Group 3: Loses 3

electrons to form 3+ ions

B3+ Al3+ Ga3+

Predicting Ionic Charges Group 4: Lose 4

electrons or gain 4 electrons?

Carbon loses 4, but the metals have

multiple charges

Predicting Ionic Charges Group 5: Gains 3

electrons to form 3- ions

N3-

P3-

As3-

Nitride

Phosphide

Arsenide

Predicting Ionic Charges Group 6: Gains 2

electrons to form 2- ions

O2-

S2-

Se2-

Oxide

Sulfide

Selenide


Group 7: Gains 1 electron to form 1- ions

F1-

Cl1-

Br1-

Fluoride

Chloride

Bromide

I1- Iodide

Predicting Ionic Charges Group 0: Stable Noble

gases do not form ions!

Predicting Ionic Charges Transition metals: Many transition elements have more than one possible charge.

Iron(II) = Fe2+ Iron(III) = Fe3+

Predicting Ionic Charges Some transition elements have only one possible charge.

Zinc = Zn2+ Silver = Ag+

Polyatomic ions

• These are groups of atoms that hang around together and have a charge.

• Hydroxide ion Ammonium ion

Here is a trick to name and know about 48 polyatomic ions.

Noble

Ga

ses

F O

ClO3

BrO3

IO3

BO3 CO3 NO3

SO4 PO4

AsO4 SeO4

TeO4 SbO4

-3 -2 -1

Now to name them…

Like the chart above

(base element name) –ate

1 more oxygen than above

Per -(base element name) – ate

1 less oxygen than above

(base element name –ite

2 less oxygen than above

Hypo –(base element name) - ite

The Octet Rule

• Atoms tend to gain, lose, or share e- until they are surrounded by 8 valence e- and are thus energetically stable.

• Exceptions do occur (and will be discussed later.)

18

Types of Bonds

1. Ionic: electrostatic attraction between oppositely charged ions (typically between a metal and a nonmetal)

2. Covalent: sharing of e- between two atoms (typically between nonmetals) molecules created

3. Metallic: “sea of e-”; bonding e- are relatively free to move throughout the 3D structure

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Covalent Metallic

Ionic

Incr

easi

ng

Dif

f. o

f EN

Ionic bonding

• Ionic bonds do not form molecules

– An ionic formula is an empirical formula (smallest whole number ratio of atoms) and doesn’t show what the structure looks like

– Ionic compounds are neutral, the positive charges balance out the negative charges.

– Ionic bonds form a lattice structure

20

Ionic Bonding • Results as atoms lose or gain e- to achieve 8

electrons in the highest shell – The bonded state is lower in energy (and therefore more

stable).

– Electrostatic attraction results from the opposite charges.

• Occurs when diff. of EN of atoms is > 1.7 (maximum is 3.3: CsF)

• Can lead to interesting crystal structures – Ionic compounds are brittle solids with high melting

points. Solids do not conduct electricity, but molten form will conduct (ions freely moving)

21

Naming Ionic Compounds Cation first, then anion

Monatomic cation = name of the element

Ca2+ = calcium ion

Monatomic anion = root + -ide

Cl- = chloride

CaCl2 = calcium chloride

If the anion is a polyatomic ion, just use the name on the polyatomic ion

CaSO4= calcium sulfate

Naming Ionic Compounds

some metal forms more than one cation

use Roman numeral in name

PbCl2

Pb2+ is cation

PbCl2 = lead(II) chloride

Metals with multiple charges (transition metals)

Writing Ionic Compound Formulas Example: Barium nitrate

1. Write the formulas for the cation and anion, including CHARGES!

Ba2+ NO3-

2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

( ) 2

Ba(NO3)2

Writing Ionic Compound Formulas Example: Iron(III) chloride


Fe3+ Cl-



Not balanced!

3

FeCl3

Writing Ionic Compound Formulas Example: Aluminum sulfide


Al3+ S2-



Not balanced!

2 3 Al2S3

Writing Ionic Compound Formulas Example: Magnesium carbonate


Mg2+ CO32-


They are balanced! MgCO3

Writing Ionic Compound Formulas Example: Zinc hydroxide


Zn2+ OH-



Not balanced!

( ) 2

Zn(OH)2

Chapter 4.2

29

Covalent Bonds

Covalent bonds form molecules

The formula is not always empirical but shows what the molecule looks like

A molecular formula shows what the molecule actually looks like

Molecular formula: C6H6 empirical: CH

30

Molecules- 2 or more atoms bound together that act as a

single, distinct object

• Theses molecules are shown in ball and stick form.

• These are also represented in structural formulas like this:

H – O H – N – H H | | | H H H – C – H | H

Condensed structural formula

• Ball and Stick Structural Condensed

Diatomics, H N F O I Cl Br

• Elements can exist as molecules.

• These elements can exist as shown.

• Notice that P and S have been added

Covalent Bonding • Atoms share e- to achieve noble gas

configuration that is lower in energy (and therefore more stable).

– Polar covalent: (different elements)

e- pulled closer to more EN atom and are shared unequally

-Nonpolar covalent: (same elements)

e- shared equally

34

Polarity: Polarity in covalent bonds

1. Polar bond: bonds in which electron density is unsymmetrical

2. Dipole: Contains a + and a - pole

3. Nonpolar bond: two atoms joined are identical.

4. Covalent bonds between two unlike atoms are always polar.

5. The extent of the polarity of a covalent bond is related to the difference in the electronegativity.

• H2 nonpolar; the hydrogens share the electrons equally

• HF polar: fluorine pulls the electrons closer so they share the electrons unequally

• In a polar molecule, one end is partially positive and one is partially negative (Dipole)

s+ s- + H—F or H—F (vector points to neg.

end)

36

Covalent bonds

• a line between atoms shows that 2 electrons are being shared H—F (single bond)

• Multiple bonds – A double line shows that 4 electrons are being shared

• O=O (double bond)

– A triple line shows that 6 electrons are being shared • N=N (triple bond) Bond length : triple < double < single Bond energy : triple>double>single

37

Bond Order • An indication of bond strength and bond length

– Single bond: 1 pair of e- shared

Ex: F2

38

•• ••

:F-F: •• ••

O=O

:N ≡ N:

Longest, weakest

Shortest, strongest

Double bond: 2 pairs of e- shared

Ex: O2

Triple bond: 3 pairs of e- shared

Ex: N2

Lewis structures show how valence electrons are arranged among atoms in a molecule.

Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron arrangemnent.

Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

Lewis Structures

Lewis structures • Central atom –least electronegative atom (carbon is always

in the center if its in the molecule)

• Hydrogen can only have a single bond because it only needs 2 electrons (can never be a central atom)

• Count the number of valence electrons in the molecule and make sure that many are in the Lewis structure. If there is a charge, add e- (if an anion) or subtract e- (if a cation).

• Complete the octet for all atoms (remember hydrogen only

needs 2 electrons)

• Another exception to the rule: Boron only wants 6 electrons around it

C H

H

H

Cl

..

..

..

Completing a Lewis Structure -CH3CCl

Add up available valence electrons:

C = 4, H = (3)(1),Cl = 7 Total = 14

Join peripheral atoms to the central atom with electron pairs.

Complete octets on atoms other than hydrogen with remaining electrons

Make carbon the central atom (least electronegative atom)

..

.. ..

Lone pairs

• Lone pairs- electrons that are not shared but belong to a single atom in a Lewis structure

• If there are too many electrons when drawing a Lewis structure, try multiple bonds

Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule.

These are resonance structures. The actual structure is an average of the resonance structures.

Bond Order & Resonance Structures

• Bond order: single bond = 1, double bond=2, triple bond = 3

• To determine bond order with resonance structures:

– Pick one bond and add up the integer bond order in one resonance structure to the same bond position in all other resonance structures.

– Divide the sum by the number of resonance structures to find bond order.

44

• Which has shorter bonds? What is the bond order in each?

– SO3 or SO32-

• Answer: SO3

Bond order for SO3 is 1 1/3

bond order of SO3 2- is 1

45

Molecular shapes: V.S.E.P.R. model

Valence-shell electron-pair repulsion theory – Because e- pairs repel, molecular shape adjusts so the valence e-

pairs are as far apart as possible around the central atom. – Electron domains: areas of valence e- density around the central

atom; result in different molecular shapes • Includes bonding e- pairs and nonbonding e- pairs • A single, double, or triple bond counts as one domain • A lone pair counts as one domain

Summary of LmABn :

L = lone or non-bonding pairs

A = central atom

B = bonded atoms

basic domains

• Linear bond angle = 180

• Trigonal planar bond angle = 120

• Tetrahedral bond angle = 109.5

• If there are only two atoms in a molecule, the geometry is linear

# of e-

domains

Molecular

geometry

formula

Predicted

bond

angle(s)

Example

(Lewis

structure with

molecular

shape)

2

Linear AB2

180º

CO2

A

|X

X

3

Trigonal

planar

AB3

Trigonal planar

120

NO3 -

LAB2

Bent

< 120º

NO21-

A

|X

X X

A |B

B B

A

|B

: B

4

or

Tetrahedral

AB4

Tetrahedral

109.5º

CH4

LAB3

Trigonal

pyramidal

< 109.5º

Ex: NH3 =

107º

NH3

L2AB2 Bent

<<109.5º

Ex: H2O =

104.5º

H2O

X

A

X

X X

X

A

X

X X

B

A

B

B B

:

A

B

B B

:

A

B

B :

Molecular Polarity

• A molecule is polar if its centers of (+) and (-) charge do not coincide. – A bond’s polarity is determined by the difference of EN between

atoms in bond. – Partial (+) and partial (-) on atoms in a polar bond can be

represented as d+ and d-.

H-Cl:

: :

H-Cl:

: :

d+ d-

Bond polarity is most often represented by an arrow that points toward the d-

(most EN atom), showing the shift in e-

density.

The sum of the bond’s dipoles in a molecule determines the overall polarity of the molecule.

Shapes whose dipoles cancel out (nonpolar) as long as all outside atoms are the same:

Linear, trigonal planar, tetrahedral- no lone pairs (if one of the outside atoms are different, the molecule is

polar)

Shapes whose dipoles don’t cancel out (polar):

Bent and trigonal pyramidal – lone pairs on central atom

Ex: Draw Lewis structures and name the molecular geometry. Is it polar or nonpolar?

CO2 H2S

Linear(nonpolar) bent (polar)

CCl4 SO2

Tetrahedral (nonpolar) bent (polar)

Intermolecular forces

Chapter 4.3

Intermolecular Forces

Forces that attract molecules to other molecules. These include:

Hydrogen bonding

Dipole-dipole attraction

van der waals forces

Melting points and boiling points reflect the strength of bonds or intermolecular forces.

Intermolecular forces are much weaker than chemical bonds

Relative Magnitudes of Forces

The types of bonding forces vary in their strength as measured by average bond energy.

Covalent bonds (400 kcal)

Hydrogen bonding (12-16 kcal )

Dipole-dipole interactions (2-0.5 kcal)

Van der waals (less than 1 kcal)

Strongest

Weakest

Polarity

A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole.

H F

d+ d-

Dipole-Dipole Attraction

Attraction between oppositely charged regions of neighboring molecules.

As the polarity increases, the attraction increases. Dipole-dipole occurs in polar molecules

Hydrogen Bonding

attraction between hydrogen and more electronegative neighboring atoms such as fluorine, oxygen and nitrogen. (very polar bonds) This is just a strong dipole-dipole attraction Hydrogen bonding occurs in polar molecules where H is bonded to F, O, or N

Hydrogen Bonding in Water

Van der waals forces

Induced or instantaneous dipoles Occurs in all molecules and atoms These are the only forces of attraction between completely nonpolar

molecules

Larger molecules (more electrons) create stronger van der waals forces

Van der waals in Hydrocarbons

List the following in order of increasing boiling points: BaCl2, H2, CO, HF, Ne

Answer: H2, Ne, CO, HF, BaCl2

nonpolar, nonpolar, polar, polar, ionic

Van der waals, van der waals, dipole-dipole, hydrogen, ionic

Chapter 4.4

Metallic bonding • Metallic elements have low I.E.; this means valence e- are

held “loosely”.

• A metallic bond forms between metal atoms because of the movement of valence e- from atom to atom to atom in a “sea of electrons”. The metal thus consists of cations held together by negatively-charged e- "glue.“

66

This results in excellent thermal & electrical conductivity, ductility, and malleability.

A combination of 2 metals is called an alloy.

Free e- move rapidly in response to electric fields, thus metals are excellent conductors of electricity.

Free e- transmit kinetic energy rapidly, thus metals are excellent conductors of heat.

Layers of metal atoms are difficult to pull apart because of the movement of valence e-, so metals are durable.

However, individual atoms are held loosely to other atoms, so atoms slip easily past one another, so metals are ductile.

Review of bonding

• Ionic compound (metal/nonmetal) creates a lattice

– Formula doesn’t tell the exact number of atoms, but instead the smallest whole number ratio of atoms (no lewis structures)

• Covalent compound (nonmetal/nonmetal) creates molecules

– Formula tells you the exact number of atoms in the molecule. (lewis structures)

IB Chemistry · Ions •When an atom gains or loses an electron, it is called an ion. (The valence...

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Transcript of IB Chemistry · Ions •When an atom gains or loses an electron, it is called an ion. (The valence...