Ionic/Covalent Models

Define lattice enthalpy

Value that shows how strong ionic bonds are in an ionic lattice. It the measure of strength of the ionic bond in different compounds influenced by ion radius and charge. The enthalpy change when one mole of a solid ionic compound is broken down into gaseous ions. It is endothermic with a postiive enthalpy change.

What is the trend of boiling points and molar mass and the exceptions?

Boiling points usually increases as molar mass increases except for NH₃ HF and H₂O . Water would be gas not liquid at room temperature if there were no hydrogen bonded.

Outline the physical properties of covalent compounds as a result of its intermolecular forces

Define dipole induced dipole

Part of van Der Watts forces. Arises when a polar molecule (permament dipole) induces a temporary dipole in a nonpolar molecule because the postitive or negative end of the polar molecule can distort the electron distribution in the nonpolar molecule, creating a temporary uneven charge distribution (dipole). this newly induced dipole is then attracted to the opposite charge of the original polar molecule, resulting in a weak attraction. Slightly stronger than london dispersion forces.

Define van Der Waals forces

Attraction between netural particles mainly due to polarity. Includes dipole-dipole forces, dipole-induced dipole forces, and London (dispersion) forces. Charged molecules are not mainly due to van Der Waals force.

Define intermolecular forces (IMF)

Attractive and repulsive forces between molecules of substance

Differences between inter and intramolecular forces

Intermolecular: Between 2 different atoms in 2 different molecules

Intramolecular: Between different atoms in same molecule

Which one has a net dipole moment and why? Cis isomer or trans isomer

Cis isomer has a net dipole moment because both electronegative atom are in the same side of the molecule while in the trans isomer the dipoles cancel out because it is diagonal.

Define lone pairs

Electrons that are not in bond

What number should electronegativity difference be to indicate that it is an ionic bond?

Over 1.7

Define octet rule

Tendency of atoms to prefer 8 valence electrons (to reach stability)

Define bond dipole

Seperation of partial postive and neg charges when electrons are unequally shared in covalent bond.

Why is that in Bohr’s Model that electrons are added in clockwise/count-clockwise direction

Orbitals. Electrons with same spin can not exist in same orbital

Define net dipole moment

Overall measure of molecule’s polarity by adding the individual bond dipole moments within molecule. If it cancels out = net dipole moment is zero (non-polar).

Detail the factors that affect lattic enthalpy

• Ionic radius- bigger the radius = decrease of electrostatic attraction between nucleus and outermost electrons (shielding effect where inner shell electrons reduce the attractive force of the nucleus on the valence electrons as they have same charge)

• Ionic charge- Higher charge = stronger attraction between ions, resulting in greater lattice enthalpy.

What classifies bonds as polar/non-polar covalent bonds based on difference in electronegativity

0: Non-polar (pure) covalent bond. Ex: Cl-Cl

0.1-0.4: Non-polar (weakly polar) covalent bond. Ex: C-H

0.5-1.7: Polar covalent bond. Ex: C-F

What factors affect polarity?

• Presence of polar bonds

• Geometry of molecule

• Polar molecules have a net dipole moment

What does it mean when an atom is partially negative?

Higher electronegativty = partial negative charge.

List the difference in electronegativty from least to greatest

Non polar bonds, polar bonds, ionic bonds

What is the effect of ionic charge on melting point

Greater the charge, the stronger the electrostatic attraction between oppositely charged ions = higher melting point.

What is the effect of ionic radius on melting point

Greater the ionic radius, the weaker the electrostatic attraction between oppositely charged ions (shielding effect) and lower the melting point.

Define lattice

A continuous, 3-D network of repeating units of positive and negative ions. The arrangement of ions in a lattice depends on the size and charge ratio of the ions.

Define volatility

The ability for a substance to change from a state of liquid to gas.

Outline the properties of ionic lattices

It has high boiling point/melting point and low volatility because the energy required to break the bonds are high.
Ionic lattice do not conduct electricity in solid state (strong electrostatic attraction) but if it is molten ions are free to move and can conduct electricity.
The ionic compound solubility depends on attraction between ions and interaction with solvent molecules.
Larger radii decreases attraction because of its electrons are further away from the nucleus.

Define polar and non polar

Polarity is the difference between electronegativity (tendency to attract electrons into a bond). Non polar means that both sides have similar electronegativity.

Describe why the shape of water has V shape

Tetrahedron structure that can’t see the other two sides due to lone pair. Valence shell electron theory (VSEPR)

Define electronegativity (χ)

Measure of ability of an atom to attract a pair of covalently bonded electrons.

Explain the trend in lattice enthalpy

Lattice enthalpy increases with ions with larger charges and smaller sizes as the energy required to overcome electrostatic forces increases.

Formula for ammonium

NH₄⁺

Formula for hydroxide

OH^-

Formula for nitrate

NO₃^-

Formula for nitrite

NO₂^-

Formula for sulfate

SO₄^2-

Formula for sulfite

SO₃^2-

Formula for phosphate

PO₄^3-

Formula for carbonate

CO₃^2-

Formula for peroxide

O₂^2-

Define coordinate covalent bond

One atom contributes to both bonding electrons to the bond instead electrons being shared. It is identical to a regular covalent bond. For example H₃O⁺ (hydronium ion), NH₄+ (ammonium ion),

Define VSEPR thoery

Valence shell electron pair repulsion theory is used to predict the geometric shape of molecules.

Define electron domain

Electrom domain is used to refer to bonds/lone pairs (valence electrons tht are not shared in covalent bond) of electrons around an atom in a molecule.

Single, double, triple bonds and lone pairs of electrons count as one electron domain.

List the electrons pairs (bond/lone pairs) that repel eachother from greatest to the least

Lone pair + lone pair > Lone pair + bonding pair > bonding pair + bonding pair

How does the position of an element in the periodic table relate to the charge of its ions?

The number of valence electrons corresponds to a charge. If the valence electron number goes past

Why is the formation of an ionic compound from its elements a redox reaction?

Electrons are gained as electrons are shared.

Explain the physical properties of ionic compounds to include volatility, electrical conductivity and solubility

Ionic compounds will dissolve when attraction between water and ions are greater than the attraction between ions.

What experimental data demonstrate the physical properties of ionic compounds?

Melting point: Ionic solids have high melting points

Solubility: Ionic compounds usually dissolve in water but not in hexane

Conductivity: Ionic compounds in aqueous solution are good conductors

How can lattice enthalpies and the bonding continuum explain the trend in
melting points of metal chlorides across period 3?

What determines the covalent nature and properties of a substance?

Electrons are shared between non-metals in order achieve a full valence shell.

How are ionic bonds formed?

Electrons are transferred from metal to non-metal. This is because of a high difference of electronegativity (non-metal attracts for electrons). If the difference is more than 1.8 = ionic bond. The electrostatic attraction between the positively charged cation and the negatively charged anion.

Explain the relationship between the number of bonds, bond length and bond strength

Bond number increases = strength increases = length decreases

Length of bond: single > double > triple
Strength of bond: single < double < triple

How does the presence of double and triple bonds in molecules influence their reactivity?

Why do noble gases form covalent bonds less readily than other elements?

It has a full valence shell, thus it is less likely to share electrons with other elements, resulting in limited covalent bonding.

Why do ionic bonds only form between different elements while covalent bonds can form between atoms of the same element?

Ionic bonds form bonds with different electronegavitiy thus if it is the same element, there will be no difference in electronegativity.

Define coordination bond and how to identify them in compounds

Type of chemical bond in which one atom supplies both electrons for a shared pair, often found in complexes with transition metals. Same as covalent bond.

How does the valence shell electron pair repulsion (VSEPR) model help predict repulsion of electron domains around a central atom?

Define electron domains

Deduce the polar nature of a covalent bond from electronegativity values

What properties of ionic compounds might be expected in compounds with polar covalent bonding?

Define dipole moment

µ (Coulomb *m)

Measure of seperation of positive and negative electrical charges within a system/measure of system’s overall polarity.

Explain the properties (structure, electrical conductivity, thermal conductivity, appearance, phsycia/chemical, uses) of diamond

Structure:
Electrical Conductivity:
Thermal Conductivity:
Appearance:
Physical/Chemical Properties:
Uses:

Define allotropes

Same element with different bonding and structural patterns thus having different chemical/physical properties.

Define London dispersion forces and the reason why non-polar molecules are gases at room temperature

Weakest intermolecular force because of the temporary inbalance of electrons because of its frequent random movement of electrons. At some instant more of the electron cloud happens to be at one end of molecule (instantaneous/induced dipole) to create temporary polarity.

The only forces that exist between non-polar molecules (only relies on London dispersion). It is the reason why many non-polar elements/compounds are gases at room temperature becasue it has low melting/boiling points (weak bonds).

Define dipole-dipole forces

Attractive forces between polar molecules (postive end of one molecule is attracted to negative end of another). Permament dipole. Strength depends on the distnace and orientation of dipole. Dipole-dipole always stronger than London dispersion forces.

These forces cause the melting/boiling points of polar compounds to be higher than those of non-polar substances.

Define hydrogen bonding

Bond where hydrogen is electrostatically attracted to the lone pair of a neighbouring higher electronegative atom (ex: NOF)

A type of dipole-dipole interaction but much stronger becasue hydrogen exerts a strong attractive force on a lone pair in the electronegative atom of a neighbouring molecule. Causes a strong boiling point because of its strong bonds.

Explain the physical properties of covalent substances to include volatility, electrical conductivity and solubility in terms of their structure

Volatility:

Electrical conductivity: Covalent substances do not contained freely moving charged particles so they are unable to conduct electricity in either solid/liquid state.

Solubility: Dissolves in non-polar substances (usually).

Order the strengths of intermolecular forces from highest to lowest

hydrogen bonding> dipole-dipole forces > London (dispersion) forces

What are some ionic compounds between only nonmetals?

NH₄ Cl (Ammonium chloride). Has ionic bonding between NH₄ ⁺ ion and chloride ion although NH₄ ⁺ is itself has covalent bonding within (N and H).

CaCO₃ (calcium cabonate): Ca²⁺ and CO₃^- ions are held together by ionic bonds although in carbonate ion ( CO₃^- ) the carbon and oxygen atoms are connected by covalent bonds.

How to determine if a molecule is polar or non polar based on molecular structure?

Symettrical VSEPR means that bond polarities cancel = non-polar.

If it doesn’t cancel out (not symmetrical)