Applying your knowledge and understanding of electronegativity (Topic 3), explain why metals and non- metals react to form ionic compounds.

Applying your knowledge and understanding of electronegativity (Topic 3), explain why metals and non-

metals react to form ionic compounds.

High difference in EN values

• Metals – low EN, weaker attraction to electrons

• Non-metals – high EN, stronger attraction to electrons

• Very different tendencies to lose or gain electrons.

EN difference of 1.8 or higher

• General rule, when the difference in electronegativity values of two elements is 1.8 or higher, the compound they formed is predominantly ionic.

Which fluoride is the most ionic?

1. NaF2. CsF3. MgF2

4. BaF2

Bonding experience!

• Take a piece of paper and draw a line down the middle• Place 20 Skittles in the center• Play 10 rounds of Rock-paper-scissors• Winner of each round moves two Skittles to his/her

side of the paper• Count the number of Skittles on either side of the

paper at the end of 10 rounds• Report result to Ms Tsui • Don’t eat your Skittles just yet!

Results

• Ms Tsui 6, Sakura 14• Radina 14, Sarah 6• Klaudia 10, Vera 10• Michael 12, August 8• Fintan 8, Isabel 12• Leo 12, Dominik 8• Chang 12, Nick 8

How can we relate this to chemistry in terms of…

• Electronegativity?• Nature of bonding?

Bonding continuum

1. Determine the position of your bond on the bonding continuum

2. “Covalent” – “Co” = together– “Valent” = valence electrons

Limitation of this activity

• In reality, when atoms share electrons to form covalent bond, they all begin with different number of valence electrons.

• Atoms share valence electrons to achieve octet structure, meaning to form a stable arrangement of eight electrons in their outer shell.

4.2 Covalent bonding4.3 Covalent structures

When electrons are shared

Essential ideas

• Covalent compounds form by the sharing of electrons.

• Lewis (electron dot) structures show the electron domains in the valence shell and are used to predict molecular shape.

4.2 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.

4.2 Applications and skills

• Deduction of the polar nature of a covalent bond from electronegativity values.

4.3 Understandings

1. Lewis (electron dot) structures show all the valence electrons in a covalently bonded species.

2. The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.

3. Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons.

4. Resonance structures occur when there is more than one possible position for a double bond in a molecule.

5. Shapes of species are determined by the repulsion of electron pairs according to VSEPR theory.

6. Carbon and silicon form giant covalent/network covalent structures.

4.3 Applications and skills1. Deduction of Lewis (electron dot) structure of molecules and ions

showing all valence electrons for up to four electron pairs on each atom. 2. The use of VSEPR theory to predict the electron domain geometry and

the molecular geometry for species with two, three and four electron domains.

3. Prediction of bond angles from molecular geometry and presence of non- bonding pairs of electrons.

4. Prediction of molecular polarity from bond polarity and molecular geometry.

5. Deduction of resonance structures, examples include but are not limited to C6H6, CO3

2- and O3.

6. Explanation of the properties of giant covalent compounds in terms of their structures.

• First, look at the structure of one hydrogen atom

• The positively charged nucleus and the negatively charged electron attract each other.

What are these bonds between atoms like?

We’ll use the most simple example, H2, to explain

• Then two hydrogen atoms approach each other . . .• . . . the electron of one atom and the nucleus of the

other atom start to attract each other.

• When the atoms are close, the attractions are so strong that a molecule of hydrogen (H2) is formed.

• The atoms are held together by the electrostatic attractions between the two nuclei and the shared pair of electrons.

• This is a single covalent bond.

Definitions• A covalent bond is the electrostatic attraction between a shared

pair of electrons and positively charged nuclei

• A double covalent bond consists of 2 pairs of shared electrons• A triple covalent bond consists of 3 pairs of shared electrons

• Simple covalent structures are small molecules held together by covalent bonds e.g. O2, N2, CO2, HCN, C2H4 (ethene) and C2H2 (ethyne)

O=O N≡N O=C=O H-C≡N H-C≡C-H

Lewis Structures

• Diagrams that show all the valence electrons in a molecule are Lewis structures

• Dot-cross diagrams show origins of shared electrons

• Non-bonding pairs of electrons are called lone-pairs e.g. water

Drawing Lewis structure instruction

• https://www.youtube.com/watch?v=nw3xVVmEAU8

Worksheet, worksheet, worksheet..

• Practice drawing Lewis structures

Naming covalent compounds

• Different rules from naming ionic compounds

CARBON DIOXIDE

Chemical formula?

CARBON DIOXIDE

CO2

What connection can you make between the name “carbon dioxide” and the chemical formula CO2

CARBON DIOXIDE

CO2

NITROGEN DIOXIDE

Chemical formula?

NITROGEN DIOXIDENO2

SULPHUR DIOXIDE

Chemical formula?

SULPHUR DIOXIDESO2

CARBON MONOXIDE

CO

SULPHUR TRIOXIDESO3

Between Nitrogen and Oxyen…

Nitrogen monoxideNitrogen dioxideDinitrogen monoxideDinitrogen trioxideDinitrogen tetroxideDinitrogen pentoxide

NONO2

N2O

N2O3

N2O4

N2O5N2O3

N2O4

N2O5

Covalent compounds and their naming

• Two or more nonmetals bonded together• Many can combine in more than one way• Prefixes are used to indicate the ratio in which

nonmetal atoms are combined together

Naming covalent compounds

• The least electronegative element is named first

• “Mono” is never used on the first element. If there is only one atom of the first element, this is shown by the absence of a prefix

• The ending of the last element named is replaced with the suffix “-ide”

Naming covalent compounds

• With oxygen, if the prefix ends in an “a” or “o”, the vowel of the prefix is dropped.– E.g. • carbon pentoxide ✔• Carbon pentaoxide ✖

Practice examples

1. CO2

2. CO3. CCl4

4. N2O5

5. SF6

6. CS2

7. Diphosphorus trinitride8. Oxygen difluoride9. Silicon tetrabromide10. Phosphorus pentachloride11. Iodine trichloride12. Dinitrogen tetroxide

Practice examples

1. CO2

2. CO3. CCl4

4. N2O5

5. SF6

6. CS2

Carbon dioxideCarbon monoxideCarbon tetrachlorideDinitrogen pentoxideSulphur hexafluorideCarbon disulphide

Practice examples

7. Diphosphorus trinitride8. Oxygen difluoride9. Silicon tetrabromide10. Phosphorus pentachloride11. Iodine trichloride12. Dinitrogen tetroxide

P2N3OF2SiBr4PCl5ICl3N2O4

Coordinate or dative bonds

• Sometimes, both the shared electrons come from the same atom.

• A coordinate or dative bond forms when a lone-pair of electrons, donates its electrons to other atoms or ions that are deficient in electrons.

Carbon Monoxide

Coordinate or dative bonds

Coordinate or dative bonds

A beaker of water contains H2O and also H+ ions (protons), and OH- ions and H3O + hydronium ions

• The charge is now delocalised across the polyatomic ion and is shown by the square brackets [ ]

The 3 examples of coordinate bond you need to know

Incomplete octets

• Some atoms, like beryllium Be and boron B, might form stable compounds with incomplete octets of electrons

1. Write the formula2. Name this molecule

Beryllium chloride

• NOT Beryllium dichloride

• Follows ionic naming rule• But draw as a covalent

molecule