Ionic/Covalent/Metallic 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.

Explain why the melting point increases going down a group

More valence shells added = more electrons = more van der Waals forces = more Energy needed to break

List the properties of phosphorus (IV) oxide

metallic/ionic, macromolecule,

List the properties of magnesium hydroxide

Physical/chemical differences of alloys and metals

Alloys changes conductivity and chemical properties like rusting. They are more resistant than metals because of its non-directional bonding, It also has different magnetic properties.

Define alloys

Mixture of two elements often a metal and another element in solid phase.

Why are alloys harder than their contsitutent metal elements alone?

Alloys have different structure/packing of cations in the lattice which is made possible due to its non-directional nature of the delocalised electrons and the fact that the lattice can accomadate ions of different sizes. This means that the layers can no longer slide over one another which decreases its malleability but increases its hardness

List the properties of silicon dioxide

Giant covalent structure made of repeating silicon and oxygen atoms.
Bond angle around 144 degrees
It is insoluble in water

High melting/boiling point because Si atoms are held in tight covalent bonds.

Every Si bonded to 4 O (1 for each lone pair = 2 in lone pair)

Brittle
No delocalised electrons (poor conductors of energy)
Conducts electricity in molten stage due to free-moving ions

Formula for cynaide

CN^-

Why do electrons want to delocalise or not?

Atom with low electronegativity mean that it is less likely to attract electrons into a bond, thus electrons are delocalised (share with every atom).

Atom with high electronegativity mean that it is more likely to attract electrons into a bond

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

structure: Each C atom is covalently bonded to 3 other C, forming hexagons in parallel layers. Each have 120˚ bond angles. Remaining valence electron = delocalised

Electrical conductivity: Good electircal conductor because of delocalised electron that gives it mobility

Thermal conductivity: Good conductor within layer of carbon but not good in other directions (perpendicular) because layers are held together with weak van Der Waals forces.

Appearance: Grey, crystalline solid

Physcial/chemical properties: Soft and slippery (layers slide over each other), brittle, high melting point, most stable allotrope.

Uses: Pencils, electrode rods in electrolysis

List physical properties vs chemical properties

Physical:

• Ionization energy (gas only)

• Melitng/boiling point

• Malleability

• Ductility

• Electrical conductivity

• State

• Density
  
  
  Chemical:

• Electron affinity (gas only. Energy change when electron is added)

• Electronegativity

• Metallic character (tendency to lose electron and form cations)

• Radioactivity

• Reactivity

• Flammability

• Acidity/Basicity

Define a metallic model

An array of positive ions thorugh which delocalized electrons can move.

Why does magneisum have a lower melting point than the elements below its group?

It has a different metalllic crystal structure that has lower amounts of delocaslised electrons = lower melting point

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

structure: Covalent solid. C atom is covalently bonded to 3 other C atoms. Forms bond angles of 120˚. Exist as single layer only (honey-comb structure). Remaining valence of C is delocalised.

electrical conductivity: Very good. One delocalised per atom = electron mobility across layers.

thermal conductivity: Best thermal conductivity. Very good.

appearance: Transparent

phsycia/chemical: Thin material. Very strong, flexible, high melting point

uses: photo-voltaic cells, touch screens, electronic devices, etc.

Explain the properties (structure, electrical conductivity, thermal conductivity, appearance, phsycia/chemical, uses) of fullorene (C₆₀)

structure: Simple covalent/molecular substance. Fixed formula (C₆₀) so it is not a covalent network solid. 20 hexagons and 12 pentagons arranged in sphere/soccer ball. Carbon is bonded to 3 others.

electrical conductivity: Poor conductors despite the fact it has delocalised electron. Little electron movement between molecules

thermal conductivity: Low thermal conductivity becasue of lack of movement of delocalised electrons

appearance: Black powder

phsycia/chemical: Light and strong. Low melting point. Reacts with K for superconducting cyrstalline material

uses: Medical, industrial devices for binding and targeting specific molecules.

What is the effect of lone pairs in bond angles of VSEPR?

Bond angle decreases as lone pairs are more repulsive than bonding pairs, forcing the bonded atoms to be closer together (smaller bond angle). This is because lone pairs are localized on one atom (not shared between atoms) thus they occupt more space around the central atom.

Distinguish between molecular and electron geometry

Molecular geometry: Arragment of atoms and only considers bonding electron domains

Electron geometry: Arrangment of elecgron domains around central atoms (includes both bonding/non-bonding electron domains)

List the properties of metals

• Sea of delocalised electrons

• Malleable (be hammered and still hold shape) because of non-directional bonding (layers of atoms can slide over eachother without breaking metallic bond)

• Ductile (drawn into wires) because metallic bonds are still intact even when structure is distorted

• High melting points

• Good conductors of heat because of its sea of delocalised electrons

• Shiny because delocalised electrons in metallic structure reflect light

• Lattice structure of metal ions, positive ions, and delocalised electrons

Define delocalised electrons and sea of electrons

Delocalised electrons: Electrons that are not bonded to a single atom or covalent bond, moving freely to bond with multiple different atoms and compounds

Sea of electrons: Model that describes how electrons behave in metal. Since electrons are delocalised, they are free to move throughout the metal structure. It is what gives metals their characteristic properties

Define metallic bond

• Non-directional= Force of attraction occurs in all directions (all the same)

• Electrostatic attraction between lattice of cations and sea of delocalised electrons

• Strength is determined by charge on metal ion and ionic radius because higher charge + smaller radius = higher melting point because attraction between outermost electron in stronger in smaller radius and more valence shells = stronger attraction between nucleus

• (Nothing on ionization energy because it is for gas phase only)

• Good electrical conductivity/thermal conductivity beause it has free moving charges and collisions= transfer of Energy)

• Stronger than ionic bonds

What are the properties of alloys

• Metallic bonding (non-directional =attraction between atoms is same from all directions and has attractive force between metal cations and delocalised sea of electrons)

• Contain cations of different sizes (makes alloys stronger and more reistant)

• Different sizes of cations prevent layers from sliding over eachother so alloys is less malleable and ductile than pure metals. The misaligment helps maintain structure

• Smaller ions can fit in gaps and increase the attractive forces

Define covalent network

Giant network/lattice of covalently bonded atoms (strong). Higher melting/boiling point. Not conductive usually unless it has free electrons.

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.

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

Unequal sharing of electrons in covalent bond due to difference in electronegativity between atoms.

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 lattice 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.

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 from the arrangement of electron pairs around central atom that minimizes the repulsion. It considers both bonding pairs and lone pairs. The shape of the moelcule is determined by the number and type of electron pairs around central atom.

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, melting/boiling point, electrical conductivity and solubility

Melting and boiling point: High melting/boiling point.

Volatility (how easily substance vaporizes aka turn to gas/vapour): Ionic compounds generally have low vapor pressure because of its strong electrostatic attraction thus it is not volatile.

Electrical conductivity: In solid state, ions are in fixed lattice that can not freely move ions to carry an electrical charge however in liquid/aqueous form it can because ions become free and carry electrical lcharge (good conductors)

Solubility in water: Solubility depends on nature of ions and interactions between water molecules and ions. Som are highly soluble, some not.

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

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?

Increases the reactivity of molecules because it is less stable. This is because electrons in multiple bonds are more accessible and can be donated to other atoms or molecules more readily.

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?

It assumes that electron pairs will arrange themselves to minimize repulsion to create a geometric shape and bond angles. It predicts the 3D arrangement of atoms in a molecule

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: C atom covalently bonded to 4 others. Tetrahedrally arranged with repetitive pattern. Bond angle 109.5˚
Electrical Conductivity: Non-conductor of electricity as all electrons are bonded = no delocalised electrons
Thermal Conductivity: Very efficient thermal conductors, better than metals because it does not have delocliased electrons = vibration between covalent bonds
Appearance: Transparent, cyrstal
Physical/Chemical Properties: Hardest known substance. Brittle and high melting point (covalent bonds)
Uses: Jewellery and ornaments and tools

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: Low melting/boiling ponits (usually) because of intermolecular forces between covalent molecules are weak compared to covalent intramolecular forces. Weaker force = less energy required to change state

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 like solvent. Polar in polar, non polar in non polar.

Order the strengths of intermolecular forces from highest to lowest

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

What are some ionic bonds 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)