1 bonding: part 1 topic 4 ionic bonding covalent bonding

BONDING: Part 1

Topic 4

•Ionic Bonding

•Covalent bonding

Let’s Review• Valence electrons

– electrons in the highest occupied energy level – always in the s and p orbitals

• normally just a draw a circle to represent these two orbitals

– determines the chemical properties of an element– usually the only electrons used in chemical bonds

Let’s Review• when forming compounds, atoms tend to achieve

the electron configuration of a noble gas (ns2np6)– highest energy level will be filled with 8 electrons the

easiest way possible– atoms of metallic elements (groups 1,2,3) lose

electrons producing cations (positive ions)• Ca becomes Ca2+

– atoms of nonmetallic elements (groups 5,6,7) gain electrons producing anions (negative ions)

• Cl becomes Cl1-

– group 4 can go either way

NeNeNNNaNa FF

MgMg2+

Cations

Anions

NN3- FF1-

...etc.

As it turns out, atoms bond together for a very simple reason: atoms like

to have full valence shells.

1+ 2+ 3-3+ 4+/- 2- 1- 0

IB Topic 4: Bonding

Only one group of elements are stable (nonreactive). What is unique about their electron structure?

Filled s & p sublevels

All other elements react in order to achieve this stable electron configuration.

Ionic Bond: Transfer of electrons; metal + nonmetalCovalent Bond: Sharing of electrons; nonmetal +

nonmetal

IONIC BONDINGIdentification

4.1.1 Describe the ionic bond as the electrostatic attraction between

oppositely charged particles.

Ionic bonding occurs when one or more electrons are transferred from the outer shell of another

When one atom gives up its electron(s), it becomes a positively charged cation.

When the other atom gains the electron(s), it becomes a negatively charged anion.

OPPOSITES ATTRACT…The ionic bond formed between the two elements is

caused by the attractive force (electrostatic attraction) between the positive and negative

Get it??

Na “gives” Cl one electron and now both atoms have a full valence shell (electron configuration of a noble gas)

4.1.2 Making ionic bonds

•Sodium gives up an electron to chlorine.

•Sodium now has a full 2nd shell, & chlorine has a full 3rd shell.

4.1.2 Describe how ions can be formed as a result of electron transfer.

Reaction between Sodium and Chlorine

• Since opposite charges attract, the Na+ and Cl- ions form an ionic bond.

• Formula NaCl (1:1 ratio)

• Name: Sodium chloride

Ions are formed when atoms gain or lose electrons

• Sodium configuration: 1s22s22p63s1 or

__ __ __ __ __ __1s 2s 2p 3s

• If Na loses one electron then it would end in 2s22p6 and be stable.

• Then sodium has 11 protons (11+), but only 10 electrons (10-) so it acquires a charge of 1+ and becomes the sodium ion, Na+. In order for it to lose an electron, something has to gain an electron

• Chlorine configuration: 1s22s22p63s23p5 or

__ __ __ __ __ __ __ __ __1s 2s 2p 3s 3p

• If Cl gains one electron then it would end in 3s23p6 and be stable.

• Then chlorine has 17 protons (17+), and 18 electrons (18-) so it acquires a charge of 1- and becomes the chloride ion, Cl-. It will gain the electron from the sodium.

Li + F Li+ F -

The Ionic Bond

1s22s1 1s22s22p5 1s2 1s22s22p6

[He] [Ne]

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -

Reaction between Magnesium and Chlorine

• Magnesium configuration: 1s22s22p63s2 or

__ __ __ __ __ __1s 2s 2p 3s

• If Mg loses two electrons then it would end in 2s22p6 and be stable.

• Then magnesium has 12 protons (12+), but only 10 electrons (10-) so it acquires a charge of 2+ and becomes the magnesium ion, Mg2+. In order for it to lose two electrons, something has to gain two electrons

Reaction between Magnesium and Chlorine

• Chlorine configuration: 1s22s22p63s23p5 or

__ __ __ __ __ __ __ __ __1s 2s 2p 3s 3p

• If Cl gains one electron then it would end in 3s23p6 and be stable.

• Then chlorine has 17 protons (17+), and 18 electrons (18-) so it acquires a charge of 1- and becomes the chloride ion, Cl-.

• Since chlorine can only gain one electron and magnesium gives up two electrons, magnesium requires two chlorine atoms.

Reaction between Magnesium and

Chlorine

• Since opposite charges attract, the Mg2+ and the 2 Cl- ions form an ionic bond.

• Determine the formula and name of the compound produced.

Reaction between Magnesium and

Chlorine

• Since opposite charges attract, the Mg2+ and the 2 Cl- ions form an ionic bond.

• Name: Magnesium chloride

• Formula MgCl2

• Magnesium gives away 2 electrons; one to each chlorine atom• Now Magnesium has a full 2nd shell and both Chlorines have a full

3rd shell

Reaction between Potassium and Oxygen

Potassium configuration: 1s22s22p63s23p64s1

• Potassium will lose 1 electron and become the potassium ion K+.

Oxygen configuration: 1s22s22p4

• Oxygen will gain 2 electrons and become the oxide ion O2-.

Two potassiums are needed to combine with one oxygen.

Formula: K2O

Name: Potassium oxide

Reaction between Aluminum and Bromine

Diagram the bonding between Al and Br,

write the formula, and give the name.

4.1.2 Describe how ions can be formed as a result of electron

transfer.

Reaction between Aluminum and Bromine

Diagram the bonding between Al and Br,

write the formula, and give the name.

Formula: AlBr3

Name: Aluminum bromide

4.1.2 Describe how ions can be formed as a result of electron

transfer.

Metal + NonmetalIonic compounds are generally composed of a metal combined

with a nonmetal(s)

Give + TakeThe metal gives it’s valence electrons to the nonmetal so it can fill

it’s valence shell.

Cation + AnionIonic compounds are the attraction between oppositely charged

Electronegativity Difference > 1.7Ionic bonds have electronegativity differences greater than 1.7

4.1.6 Predict whether a compound of two elements would be ionic from the position of the

elements in the periodic table or from their electronegativity values.

0.1 – 1.0

1.1 – 1.7

0.0 covalent, nonpolar

covalent, slightly polar

covalent, very polar

electronegativtydifference

probable type of bond

IONIC BONDINGFormulas & Naming

Writing formulas for ionic compounds

• Write symbol of cation and then anion

• Add subscripts to balance the charges– calcium bromide

• Ca2+ and Br1- is CaBr2

– potassium sulfide

• K+1 and S2- is K2S

– iron(III) oxide

• Fe+3 and O2- is Fe2O3

“Crisscross” method The ionic charge number of each ion is crossed

over and becomes the subscript for the other ion

Reduce to lowest ratio

Naming Ionic CompoundsNaming Ionic Compounds

• 1. Cation first, then anion

• 2. Monatomic cation = name of the element

• Ca2+ = calcium ion

• 3. Monatomic anion = root + -ide

• Cl = chloride

• CaCl2 = calcium chloride

Naming Ionic Compounds

Examples:

sodium chloride

zinc iodide

aluminum oxide

Learning Check

Complete the names of the following binary compounds:

Na3N sodium ________________

KBr potassium ________________

Al2O3 aluminum ________________

MgS _________________________

Names of Variable IonsNames of Variable Ions

Transition metals (except Ag, Zn, Cd) REQUIRE Roman Numerals because they can have more than one possible charge.

I II III IV V VI VII 1 2 3 4 5 6 7

FeCl3 (Fe3+) iron (III) chlorideCuCl (Cu+1) copper (I) chlorideSnF4 (Sn4+) tin (IV) fluoridePbCl2 (Pb2+) lead (II) chloride

Fe2S3 (Fe3+)iron (III) sulfide

Naming ionic compounds• Binary Compounds

– cation is written first, followed by the anion with and –ide ending• Cs2O cesium oxide

• SrF2 strontium fluoride

• CuO copper(II) oxide– oxygen is always 2- and therefore copper will be 2+

• Cu2O copper(I) oxide– oxygen is 2- and therefore needed two copper atoms

with 1+ charge

• Ionic compounds with transition metals• indicate charge after metal with Roman numerals

• The overall charge of the compound should = 0

• Ionic compounds with transition metals• indicate charge after metal with Roman numerals

• The overall charge of the compound should = 0

FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride

FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride

Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide

– SnF2 ?

• fluorine is always 1- and therefore tin will be 2+

– tin(II) fluoride

– SnS2 ?

• sulfur is always 2- and therefore tin will be 4+

– tin(IV) sulfide

Learning Check

Complete the names of the following binary

compounds with variable metal ions:

FeBr2 iron (_____) bromide

CuCl copper (_____) chloride

SnO2 ___(_____ ) ______________

Fe2O3 ________________________

SnF2 ________________________

Learning Check

Complete the names of the following binary

compounds with variable metal ions:

FeBr2 iron (II) bromide

CuCl copper (I) chloride

SnO2 tin (IV) oxide

Fe2O3 iron (III) oxide

SnF2 tin (II) fluoride

4.1.5 State that transition elements can form more than

one ion.MOST transition metals REQUIRE Roman

Numerals because they can have more than one possible charge.

BUT there are ALWAYS exceptions to rules

(in chemistry)• These transition metals only form ONE ion:

Ag+1, Zn+2 and Cd+2

Label these on your periodic table!!

NONO33--

nitrate ionnitrate ion

NONO22--

nitrite ionnitrite ion

Polyatomic IonsPolyatomic Ions

Know these!!!Know these!!!NH4 +1 Ammonium

HCO3 –1 Bicarbonate aka hydrogen carbonate

CO3 – 2 Carbonate

Cr2O7 –2 Dichromate

OH –1 Hydroxide

NO3 –1 Nitrate

MnO4 –1 Permanganate

PO4 –3 Phosphate

SO4 –2 Sulfate

Naming Ternary Compounds

Contains at least 3 elementsThere MUST be at least one polyatomic ion

(it helps to circle the ions)Examples:

NaNO3 Sodium nitrate

K2SO4 Potassium sulfate

Al(HCO3)3 Aluminum bicarbonate

Aluminum hydrogen carbonate

Learning Check

Match each set with the correct name:

1. Na2CO3

2 . Ca(HCO3)2

Ca3(PO4)2

Learning Check

Match each set with the correct name:

1. Na2CO3 a) magnesium sulfite

MgSO3 b) magnesium sulfate

MgSO4 c) sodium carbonate

2 . Ca(HCO3)2 a) calcium carbonate

CaCO3 b) calcium phosphate

Ca3(PO4)2 c) calcium bicarbonate

Ionic Nomenclature

Writing Formulas

• Overall charge must equal zero.– If charges cancel, just write symbols.– If not, use subscripts to balance charges.

• Use parentheses to show more than one of a particular polyatomic ion.

• Use Roman numerals indicate the ion’s charge when needed (transition metals)

• Don’t show charges in the final formula.

Ternary Ionic Nomenclature

Sodium SulfateNa+ and SO4 -2

Na2SO4

Iron (III) hydroxideFe+3 and OH-

Fe(OH)3

Ammonium carbonateNH4

+ and CO3 –2

(NH4)2CO3

Learning Check

1. aluminum nitrate

a) AlNO3 b) Al(NO)3 c) Al(NO3)3

2. copper(II) nitrate

a) CuNO3 b) Cu(NO3)2 c) Cu2(NO3)

3. Iron (III) hydroxide

a) FeOH b) Fe3OH c) Fe(OH)3

4. Tin(IV) hydroxide

a) Sn(OH)4 b) Sn(OH)2 c) Sn4(OH)

Formula to NameFe(NO3)3

Choose the correct name for the compound

1. Iron trinitrate

2. iron(I) nitrate

3. iron(III) nitrite

4. iron(III) nitrate

5. none of the above

next problemPolyatomic IonsPeriodic Chart

Name to Formulasodium chlorite

Choose the correct formula for the compound

1. NaCl

2. NaClO

3. NaClO2

4. Na(ClO)2

5. none of the above

next problemPrefixesPeriodic Chart

Mixed Practice!

Name the following:

1. Na2O

2. CaCO3

3. PbS2

4. Sn3N2

5. Cu3PO4

6. MgF2

Mixed Up… The Other Way

Write the formula:

1. Copper (II) chlorate

2. Calcium nitride

3. Aluminum carbonate

4. Potassium bromide

5. Barium fluoride

6. Cesium hydroxide

IONIC BONDINGPhysical Properties

Properties of Ionic Compounds

• OVERARCHING CONCEPT– Anions and cations are strongly attracted to each

other and difficult to separate an ionic compound– Fuse School video:

https://www.youtube.com/watch?v=TxHi5FtMYKk&list=PLW0gavSzhMlReKGMVfUt6YuNQsO0bqSMV&index=44

Properties of Ionic Compounds• Crystalline “lattice” structure

– repeating arrays of cations and anions

– an ionic lattice of alternating positive and negative ions • Each sodium ion is surrounded by up to 6 chloride ions and

each chloride ion is surrounded by up to 6 sodium ions.

Properties of Ionic Compounds• Hard and brittle

– When pressure is applied, ions of like charge will be forced closer to each other

– The electrostatic repulsion can split the crystal

Properties of Ionic Compounds

• High melting and boiling points – High temperatures are required to overcome the

attraction between the cations and anions

– Therefore, it takes a lot of energy to break apart the electrostatic forces allowing them to melt or boil

• Volatility: how easily a substance turns into a gas– very low as electrostatic forces between cations

and anions is very strong– How often have you seen or heard of gaseous

salt??

• Solubility: the ability to dissolve – Most will dissolve in other polar solvents such as water

– Ions dissociate, or separate from each other, when they dissolve, breaking apart the lattice structure

– ions keep their charges in solution

– Solubility of salt in water: http://www.youtube.com/watch?v=xdedxfhcpWo&feature=related

Conductivity• Electrical conductivity: the ability to allow for

the flow of electrons• Substances must possess Freely Moving

Charged Particles (ions or subatomic particles) – This occurs in…

• molten ionic compounds (+ and – ions can move)– http://www.dynamicscience.com.au/tester/solutions/che

mistry/bonding/bonding5.htm

• ionic compounds in aqueous solution (dissolved in water)

– water pulls apart + and – ions and allows them to move

– This does not occur in solid ionic compounds• ions are bound so tightly to each other that there is nowhere

for electrons to flow

4.1.8 Describe the lattice structure of ionic compounds.

http://ed.ted.com/lessons/how-atoms-bond-george-zaidan-and-charles-morton

https://www.youtube.com/watch?v=lhC42qxk5kQ

COVALENT BONDINGLet’s share

IB Topic 4: Bonding4.2: Covalent bonding

Essential Idea: Covalent compounds form by the sharing of electrons.

Nature of Science:Looking for trends and discrepancies – compounds

that contain non-metals have different properties from compounds that contain non-metals. (2.5)

Use theories to explain natural phenomena – Lewis introduced a class of compounds which share electrons. Pauling used the idea of electronegativity to explain unequal sharing of electrons (2.2)

Understandings:1. A covalent bond is formed by the electrostatic

attraction between a shared pair of electrons and the positively charged nuclei.

2. Single, double, and triple covalent bonds involve one, two, and three shared pairs of electrons, respectively.

3. Bond length decreases and bond strength increases as the number of shared electrons increases.

4. Bond polarity results from the difference in electronegativities of the bonded atoms.

Applications and Skills:1. Deduction of the polar nature of a covalent bond

from electronegativity values

4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei

Covalent bonds occur between nonmetals since both want to gain electrons.

Atoms share valence electronsthis is still in order to achieve an noble gas

electron configuration (stable and less energy)

exists where groups of atoms (or molecules) share one or more pair/s of electrons

Each hydrogen now has the electron configuration of the nearest noble gas- helium

The electrons in the bond are electrostatically attracted by both nuclei,

so that it forms a directional bond between the two atoms.

4.2.2 Describe how the covalent bond is formed as a result of

electron sharing.Covalent bonds occur between nonmetals since both want to

gain electrons. They share electrons to achieve a stable configuration

Electron Dot Diagrams

Need 1 electron: H F Cl Br I

Need 2 electrons: O S Se Te

Need 3 electrons: N P As

Need 4 electrons: C Si

4.2 U1. Electron sharing

Key Terms:Lone pairs: electrons on a dot diagram that are

already paired (also called non-bonding pairs)

Shared pairs (Bond pair): electrons that are shared in a covalent bond

H Shared pair

Lone pair

4.2.2 A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

7e- 7e-

8e- 8e-

Lewis structure of F2

lone pairslone pairs

single covalent bond

4.2 U2. Bond Types

Double covalent bonds

Sharing 2 pair of electrons

Examples: O2

Single covalent bonds

Examples:H2 H H H H

HCl H Cl H Cl

4.2 U2. Bond Types

Triple covalent bonds

Examples:N2

H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N N N

triple bondtriple bond

CCl4 - Covalent

HCl - Covalent

MgF2 - Ionic

[ F ]2– [Mg]2+

H2O - Covalent

NH3 - Covalent

NaCl - Ionic

[ Cl ]– [Na] +

OH– - Covalent

H2 - Covalent

HCl - Covalent

H Cl H Cl

CO2 - Covalent

Na2O - Ionic

[ O ]2– [Na]2+

O2 - Covalent

I2 - Covalent

[ O ]32– [Al]2

Al2O3 - Ionic

NH3 - Covalent

O3 - Covalent

Modeling Covalent Bonds• Molecular model sets: modeling covalent

bonding. Create the 8 molecules listed on the next slide and complete a data table similar to the following:

Formula (given)

Line Diagram Dot Diagram

You will be making 8 molecules!!

Modeling Covalent BondsRULES:•All bonding sites must have a bond

• Exception: nitrogen only has 3 bonding sites… the 4th should have a “hat”

•Each bond must be attached to two atoms•Molecules will not be circular

Modeling Covalent Bonds• 1 bond = 2 shared electrons• Red = oxygen• Black = carbon• White = hydrogen• Blue = nitrogen (3 bonding sites + 1 “hat”)• Green & SilverSilver = halogensgens• H2O ● H2

• C2H4 ● O2

• CO2 ● C6H12

• NH3 ● N2

4.2 U3. Bond lengths and Strength

The more pairs of Electrons that are

shared between two atoms (bonds) in a molecule will

make the attraction between

the atoms • stronger

• and shorter

Lengthnm

Strength(kj mol-1)

C-O 0.143 356

C-C 0.154 348

C=O 0.121 736

C=C 0.134 657

C C 0.120 908

So, if you compare a single bond to a triple bond, • Which one is stronger?• Which one is shorter?

Bond Strength

Triple bond > Double Bond > Single Bond

Strongest Medium Weakest

Bond Lengths

Triple bond < Double Bond < Single Bond

Shortest Medium Longest

Know difference in strength and length between single, double, and triple bonds between two

carbon atoms.

Know the bond length difference between the C and O in the carboxyl group

Let’s practice!

Draw C3H6 (or CH3CHCH2)

• Identify the longest carbon-to-carbon bond• Identify the strongest carbon-to-carbon bond

• Now try: CH3COOH

• Now try: HCOOCH3

COVALENT BONDINGBOND Polarity

Covalent bonds occur between nonmetals since both want to gain electrons.

Atoms share valence electronsthis is still in order to achieve an noble gas

electron configuration (stable and less energy)exists where groups of atoms (or molecules)

share one or more pair/s of electronsSometimes the sharing of electrons between

atoms is unequal, leading to polarity

Polarity• Covalent bonds can either be nonpolar or polar• Shared bonding electrons pairs are sometimes

pulled (as in a “tug-of-war”) between atoms– Equal sharing is non-polar– Unequal sharing is polar

4.2.6 Predict the relative polarity of bonds from electronegativity values.

Nonpolar Covalent Bond:Nonpolar bonds form when electrons are shared equally

between two atoms with the same electronegativity values

2.2 2.2 – 2.2 = 0

Nonpolar covalent bonds

• Since atoms in the bond have the same electronegativity values, they pull on the shared pair of electrons equally, so no polarity = nonpolar

• Always the case in diatomic molecules

–BrINClHOF meaning…

•Br2 I2 N2 Cl2 H2 O2 F2

Memorize these!!!

Polar Covalent Bonds• Atoms in the bond pull the shared pair of

electrons unequally since they have different electronegativities

• Results in a dipole because it has two poles use the symbol + or – for areas that are slightly

positive or negatively charged

• More electronegative atoms have a greater attraction for electrons• A number is assigned to each element to

quantify its attraction to a pair of electrons in a shared in bond (example- F is 4.0)

• Polar Covalent Bond: Polar bonds form when electrons are shared unequally between two atoms due to a difference in electronegativity

So why are some BONDS polar?

• Atoms with the higher electronegativity give that “side” of the molecule a slightly negative charge (-)

• Atoms on the “other side” with a lower electronegativity therefore have a slightly positive charge (+)

• The greater the difference in electronegativity, the more polar the bond will be

So why are some BONDS polar?

covalent, non-polar

covalent, polar

BONDING Practice

Arrange the following BONDS from most to least polar:

a) N–F O–F C–F

a) C–F, N–F, O–F

b) C–F N–O Si–F

b) Si–F, C–F, N–O

c) Cl–Cl, B–Cl, S–Cl

c) B–Cl, S–Cl, Cl–Cl

4.2 U4. Electronegativity

Electronegativity Difference

Nonpolar covalent bonds have electronegativity difference is 0

Polar covalent bonds have electronegativity differences above 0, but less than 1.7Ionic bonds have electronegativity

differences greater than 1.7http://animatedchemistry.org/?p=528

4.2.6 Predict the relative polarity of bonds from electronegativity values.

Nonpolar vs Polar Covalent Bonds:

4.2 U4. Electronegativity

Nonmetal + NonmetalCovalent compounds generally occur between

elements that are closer to each other on the periodic table (with the exception of

hydrogen)

Electronegativity DifferenceCovalent bonds have electronegativity differences

from 0 to 1.7. A bond with electronegativity difference of 1.7 is

classified as covalent, NOT ionic.

0.1 – 1.0

1.1 – 1.7

0.0 covalent, nonpolar

covalent, slightly polar

covalent, very polar

electronegativtydifference

probable type of bond

Nonpolar Covalent

share e-

Polar Covalent

partial transfer of e-

transfer e-

Increasing difference in electronegativity

4.2.6 Classification of bonds by difference in electronegativity

Polarity Practice

Let’s get some

practice!!!

COVALENT BONDINGFormulas & Naming

Covalent Compound Naming

• Covalent compounds are most often called called molecules

Molecular (Covalent) Namingfor two nonmetals

• Prefix System (binary compounds)

1. Less electronegative atom comes first.

2. Add prefixes to indicate # of atoms. Omit mono- prefix on the 1st element. Mono- is REQUIRED on the 2nd element.

3. Change the ending of the 2nd element to -ide.

PREFIXmono-di-tri-tetra-penta-hexa-hepta-octa-nona-deca-

NUMBER123456789

Molecular Naming Prefixes

Molecular (Covalent) Namingfor two nonmetals

Exceptions to the prefix method

•H2O = water

– not dihydrogen monoxide

•NH3 = ammonia

– not nitrogen trihydride

ELEMENTS THAT EXIST AS DIATOMIC MOLECULES

Remember:

BrINClHOF

These elements only exist as

PAIRS. Note that when they

combine to make compounds, they

are no longer elements so they are no longer in

pairs!

• CCl4

• N2O

• SF6

• carbon tetrachloride

• dinitrogen monoxide

• sulfur hexafluoride

Molecular Naming: Examples

• arsenic trichloride

• dinitrogen pentoxide

• tetraphosphorus decoxide

• AsCl3

• N2O5

• P4O10

Molecular Formula Examples

Practice 1

Fill in the blanks to complete the following names of covalent compounds.

CO carbon ______oxide

CO2 carbon _______________

PCl3 phosphorus _______chloride

CCl4 carbon ________chloride

N2O _____nitrogen _____oxide

NH3 ______________

watch out for that last one…

Practice 1

Fill in the blanks to complete the following names of covalent compounds.

CO carbon monoxide

CO2 carbon dioxide

PCl3 phosphorus trichloride

CCl4 carbon tetrachloride

N2O dinitrogen monoxide

Practice 2

1. P2O5 a) phosphorus oxide

b) phosphorus pentoxide

c) diphosphorus pentoxide

2. Cl2O7 a) dichlorine heptoxide

b) dichlorine oxide

c) chlorine heptoxide

3. Cl2 a) chlorine

b) dichlorine

c) dichloride

Learning Check

1. P2O5 a) phosphorus oxide

b) phosphorus pentoxide

c) diphosphorus pentoxide

2. Cl2O7 a) dichlorine heptoxide

b) dichlorine oxide

c) chlorine heptoxide

3. Cl2 a) chlorine

b) dichlorine

c) dichloride

COVALENT BONDINGPhysical Properties

Covalent• Strong intramolecular forces (between atoms)• Weak intermolecular forces (between

molecules)• Usually liquids or gases at room temp or soft

solid (http://www.rsc.org/periodic-table at 25°C)

– strength of polarity determine mp and bp• greater polarity = higher mp and bp

• Polar substances are soluble in water, but nonpolar are not

• Do not conduct electricity

Solubility of Molecules

• “Like dissolves like”– Polar substances tend to dissolve in polar

solvents , such as water– Non-polar substances do not dissolve in

polar substances like water, but do tend to dissolve in non-polar solvents

Conductivity of Molecules• Substances must possess Freely Moving

Charged Particles (ions or subatomic particles) – Since they do not contain ions, they have no means

to carry an electrical current (flow of electrons)

Type of Bonding

Melting

Boiling Point

Volatility

Electrical Conductivit

Solubility in Non-

polar Solvent

Solubility in Polar

Solvent

Ionic Bonding

high high low Yes (molten or aqueous)

No Yes (most)

Polar Covalen

high varies high No No No

Nonpolar

Covalent

High High high No (except graphite

and graphene)

Characteristics of Covalent Bonds

If the type of bonding is known, physical properties can be predicted

Type of Bonding

Melting Point

Boiling Point

Volatility

Electrical Conductivity

Solubility in Non-polar

Solvent

Solubility in

Polar Solvent

Non-polar Low Low High No Yes No

Polar No No Yes

4.5.1 Compare and explain the properties of substances resulting from different

types of bonding

If the type of bonding is known, physical properties can be predicted

Type of Bonding

Melting Point

Boiling Point

Volatility

Electrical Conductivity

Solubility in Non-polar

Solvent

Solubility in

Polar Solvent

Metallic Bonding

Yes No No

Covalent No No No

Giant Covalent

High High Low No (except graphite)

1 bonding: part 1 topic 4 ionic bonding covalent bonding

gain electrons

stable electron configuration

electron structure

atoms bond

result of electron transfer

sodium chloride

sodium ion

ionic bondssodium

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