Chemical Bonding

Electrons are transferred from one atom to another

Electrons are shared equally between two atoms

Electrons are shared unequally between two atoms

Positive ions are attracted to a sea of electrons

1

2

3

4

Metallic

Covalent

Ionic

Polar Covalent

$O_2$

$CO_2$

$F_2$

$N_2$

$O_{2}$

$N_{2}$

$F_{2}$

$Cl_{2}$

Non-polar covalent

Polar covalent

Ionic

Metallic

$CuPO_{4}$

$Cu_{2}PO_{4}$

$Cu_{3}(PO_{4})_{2}$

$Cu_{2}(PO_{4})_{3}$

16; because there are twice as many cations as anions, each cation must form twice as many interactions to maintain lattice stability.

4; because the ratio of the coordination numbers must be the inverse of the stoichiometric ratio of the ions in the formula.

8; because the coordination number is a fixed geometric property of the crystal system and is identical for both types of ions.

2; because the coordination number of the cation is directly determined by its subscript in the empirical formula.

$H_2O$

$NH_3$

$CH_4$

$CO_2$

$Zn$ (Group 12)

$Cu$ (Group 11)

$Cr$ (Group 6)

$K$ (Group 1)

Atoms want to have 8 protons

Atoms want to have 8 neutrons

Atoms want to have 8 valence electrons

Atoms want to have 8 total electrons

Nitrogen oxide

Dinitrogen tetroxide

Nitrogen tetraoxide

Dinitrogen oxide

1

2

3

0

1

2

3

4

They only occur between atoms of the same element.

They involve the sharing of protons.

They result from the attraction between shared electrons and the positive nuclei.

They are weaker than intermolecular forces.

6

8

10

12

Nitrogen

Oxygen

Carbon

Sulfur

$180^{\circ}$

$120^{\circ}$

$< 120^{\circ}$

$109.5^{\circ}$

$XY$

$XY_{2}$

$XY_{3}$

$X_{3}Y$

$HF$

$CF_{4}$

$NH_{3}$

$H_{2}O$

Pure metals are mixtures; alloys are pure

Pure metals have one type of atom; alloys have two or more

Alloys have no electrons

Pure metals do not conduct electricity

The vibration of positive ions

The reflection of light by delocalized electrons

The absorption of light by the nucleus

The movement of heat through the lattice

Electrical conductivity

High density

Malleability

High melting point

Solid

Liquid

Gas

Plasma

The fixed cations (B) are released when the metal melts, allowing them to carry the electric current instead of the electrons.

The delocalized electrons (C and F) remain free to move throughout the substance regardless of the arrangement of the cations.

The regular lattice structure of the solid metal (S) actually resists electron flow, while the liquid state (D) removes all resistance.

In the liquid state (D), the cations and electrons merge to form neutral atoms that can move more freely through the metal.

The delocalized electrons evaporate

The electrostatic attractions are overcome enough for the ions to move past each other

The positive ions turn into negative ions

The metallic bonds become covalent bonds

Magnesium has a larger atomic radius

Sodium has more delocalized electrons

Magnesium has a higher nuclear charge and more delocalized electrons

Sodium ions have a higher charge density than Magnesium ions

Metal B, because the higher charge of its cations increases the electrostatic pull on the delocalized electrons, making them more stable.

Metal A, because the larger atomic radius and lower cation charge result in a weaker attraction to the delocalized 'sea of electrons.'

Metal B, because the smaller radius of its cations creates a more compact structure that allows electrons to vibrate with higher kinetic energy.

Metal A, because the lower density of delocalized electrons reduces the probability of electron-electron collisions during the escape process.

The bonds between atoms are easily broken and not reformed.

The sea of electrons acts like a flexible 'glue' that allows the cations to slide past each other while maintaining the attraction.

The atoms in a metal are not actually bonded; they are just packed tightly.

The metallic bond is purely magnetic, allowing for easy repositioning.

Strength decreases because the atoms get heavier.

Strength increases because the charge on the cation increases and the ionic radius decreases.

Strength remains constant because the number of shells is the same.

Strength decreases because $Al$ is more non-metallic than $Na$.

2

8

56

1

1

2

3

4

4

6

16

34

Circles

Brackets

Triangles

Arrows

1

2

3

5

2

4

8

32

Mg with 1 dot

Mg with 2 dots

Mg with 7 dots

Mg with 8 dots

32

34

36

38

There is a double bond between the two Oxygen atoms

There is a single bond between the two Oxygen atoms

Each Hydrogen is bonded to both Oxygen atoms

The Oxygen atoms are not bonded to each other

0

1

2

3

They have large surface areas

A large amount of energy is needed to break many strong covalent bonds

They are found deep in the Earth

They contain very few atoms

Isotopes

Isomers

Allotropes

Polymers

Sodium chloride

Diamond

Graphite

Polyethene

Each carbon is bonded to 4 others in a 3D frame

Each carbon is bonded to 3 others in layers

Carbon atoms are not bonded to each other

Carbon atoms form double bonds with 4 others

Its ability to conduct electricity when molten.

Its giant covalent structure which provides stability and high melting points.

Its reaction with water.

Its low density.

It only has two carbon atoms in its structure.

It is only one atom thick.

It can only be seen from two sides.

It is a square shape.

They are very weak and break easily.

They have a very high length-to-diameter ratio.

They are only used as insulation.

They do not contain covalent bonds.

The melting point increases because the atoms become heavier, requiring significantly more thermal energy to disrupt the vibration of the rigid lattice.

The melting point decreases because the larger atomic radii lead to longer covalent bonds, which weakens the electrostatic attraction between the nuclei and the shared electrons.

The melting point increases because the larger electron clouds in Germanium create stronger instantaneous dipole-induced dipole forces throughout the 3D network.

The melting point remains relatively constant because all three elements maintain a coordination number of four, ensuring the same number of covalent bonds must be broken.

Due to the lone pair repulsion on the oxygen atoms and the nature of the bridging bonds.

Because the oxygen atoms are $sp^2$ hybridized.

Because the silicon atoms are in a planar arrangement.

Because the structure is actually ionic.

Tetrahedral

Trigonal Planar

Linear

Pyramidal

Carbon

Silicon

Nitrogen

Iron

Weak forces

Its rigid three-dimensional lattice of covalent bonds

Its layered structure

Delocalized electrons

Diamond is made of carbon, graphite is not

Diamond is soft, graphite is hard

Diamond does not conduct electricity, graphite does

Diamond has a low melting point, graphite has a high one

Each oxygen is shared by 2 silicon atoms

Each oxygen is shared by 4 silicon atoms

Oxygen atoms are not shared

Each silicon shares 1 oxygen

The ratio of silicon to oxygen atoms is constant throughout the lattice, but the total number of atoms depends on the crystal's size.

The substance consists of discrete $SiO_2$ molecules that are too large to be measured individually.

The covalent bonds between the silicon and oxygen atoms are temporary and change as the substance is heated.

Silicon dioxide is a mixture of separate silicon and oxygen elements rather than a single compound.

The rigid, directional covalent bonds do not allow atoms to slide past one another without the structure breaking.

The weak intermolecular forces between the carbon atoms are easily overcome by the force.

The delocalized electrons create planes of weakness that cause the crystal to split.

The carbon atoms are arranged in layers that repel each other when the atoms are pushed together.

They conduct electricity.

They contain carbon atoms.

They have high melting points.

They are soft and slippery.

Diamond > Silicon Dioxide > Ice ($H_{2}O$)

Ice ($H_{2}O$) > Silicon Dioxide > Diamond

Silicon Dioxide > Ice ($H_{2}O$) > Diamond

Diamond > Ice ($H_{2}O$) > Silicon Dioxide

It becomes an insulator.

It becomes a p-type semiconductor by creating 'holes'.

It becomes an n-type semiconductor by providing extra delocalized electrons.

The lattice collapses because the atoms are different sizes.

Covalent bonding

Ionic bonding

London dispersion forces (van der Waals)

Hydrogen bonding

Dipole-dipole force

Covalent bond

London dispersion force

Hydrogen bond

A puddle of water evaporating in the sun

Iron rusting in the presence of oxygen

An acid neutralizing a base

Electrolysis of water into Hydrogen and Oxygen

intermolecular; intramolecular

intramolecular; intermolecular

ionic; covalent

weak; strong

They are only found in polar molecules.

They result from temporary shifts in electron density.

They are the strongest type of bond.

They involve the sharing of protons.

Dipole-dipole forces occur between non-polar molecules.

London dispersion forces involve permanent dipoles.

Dipole-dipole forces involve permanent dipoles, while London forces involve temporary dipoles.

London dispersion forces are always stronger than dipole-dipole forces.

Stronger intermolecular forces lead to higher vapor pressure.

Stronger intermolecular forces lead to lower vapor pressure.

Intramolecular forces determine vapor pressure, not intermolecular forces.

There is no relationship between forces and vapor pressure.

$H_2O$

$NH_3$

$CH_4$

$CH_3OH$

Dispersion forces in $Xe$

Dipole-dipole forces in $CH_{3}Cl$

Hydrogen bonds in $HF$

Covalent bonds in $C(diamond)$

$SiO_{2}$

$P_{4}O_{10}$

$SO_{2}$

$CO_{2}$

The forces are broken entirely

The forces become stronger as molecules move faster

The molecules gain enough kinetic energy to partially overcome the attractions

The intramolecular bonds begin to vibrate more, but the intermolecular forces remain unchanged