Unit 2.1: Intramolecular Bonding Flashcards
Intramolecular bonding
bonding within molecules or between atoms
Intermolecular bonding
the attraction in between separate molecules
properties of an ionic compound
conductivity:
* can NOT conduct electricity when solid because the ions are fixed in lattice and cannot move
* conducts electricity when molten (liquid) and aqueous because the ions can now move
melting + boiling point
* very high: they have very strong charges or full positive and negative charges. with this, there is a strong electrostatic attraction between oppositely charged ions and so it takes a lot of energy to break them apart.
hardness + malleability:
* hard but brittle. they are not malleable because each ion has a very specific location. even the slightest shift can cause like charges of the ions to repel apart.
solubility:
* usually very water soluble as the partial charges of polar water molecules attract the charged ions and pull them out of the crystal lattice.
properties of covalent substances (molecular)
conductivity:
* do NOT conduct because they don’t have any free moving electrons
melting + boiling point
* relatively low as the intermolecular forces are weaker than intramolecular bonding
hardness + malleability:
* huge range
solubility:
* huge range as it depends on whether they form polar or non polar covalent bonds
properties of covalent substances (macromolecular/giant structures)
conductivity:
* can conduct electricity ( graphene and graphite) because these substances do have delocalised electrons
melting + boiling point
* very high as they only have intramolecular bonds or, in other words, no molecules
hardness + malleability:
* very hard and brittle due to rigid crystal network structure
solubility:
* insoluble
properties of metals and alloys
conductivity:
* ALWAYS conduct electricity when solid and liquid
* its also has high thermal conductivity as closely packed ions and delocalised electrons efficiently transfer thermal energy
melting + boiling point
* range, though usually lower than ionic compounds
hardness + malleability:
* malleable because metal atoms can slide past each other/metallic bonds can be broken and then reformed in the new location. the attraction from delocalised electrons remains even as the metal ions move.
solubility:
* not soluble
luster: because delocalised electrons reflect light, causing a shiny appearance
metallic bond
the electrostatic attraction between positive metal cations and delocalised electrons.
atoms have low electronegativity and valence electrons are delocalised
the metal ions are in fixed positions (when solid) and can be described as a lattice of positive ions. the valence electrons of all of the metal atoms become delocalised and are able to move throughout a lattice structure of metal ions. this is due to the low attraction that each metal has for its own valence electrons. as the electrons are in constant motion throughout the lattice, the metallic bonds exist in all directions around each metal ion. with this, metallic bonding can be described as non-directional.
covalent bond
electrostatic attraction between 2 positive nuclei and the shared electron pairs between them.
high electronegativity and the valence electrons are shared in pairs.
non metals with non metals.
either molecular (most compounds) or giant covalent structures like diamond (carbon), silicone, silicone dioxide, graphite, and graphene.
ionic bond
strong electrostatic attraction between oppositely charged ions.
atoms have low and high electronegativity and the valence electrons are transferred.
between metals (low electronegativity) and non-metals (high electronegativity)
forms a crystal lattice where ions are surrounded by 6 ions of the opposite charge.
electronegativity
the ability of an atom to attract an electron pair when bonding
ionic hydrates
many ionic compounds contain water molecules as part of their crystalline structure. there are known as ionic hydrates.
water molecules (with partial charges) are attracted to the ions (with full charges) through ion-dipole interactions, filling gaps in the lattice of ions. these attractions are strong enough that water from the air will get attracted into many ionic compounds, hydrating them in the process.
by heating the ionic compound, water can be driven out of the ionic lattice, leaving behind the anhydrous form of the ionic compound (no water).
ionic compounds (number of particles)
ionic compounds are not molecules. they exist as huge crystal lattices of ions. the chemical formula of an ionic compound therefore represents the simplest ratio of the ions in the crystal. the chemical formula of an ionic compound can be referred to as a formula unit.
** iB loves to ask about ionic compounds and ionic hydrates in mole questions where you need to determine the number of particles. this requires extra careful reading as determining the number of atoms or the number of ions can be slightly different depending on if polyatomic ions are involved.
formula unit
the term used when referring to the number of particles of an ionic compound, since saying molecules is inaccurate.
polyatomic ions
Zn^2- Zinc Cation
NH4^+ Ammonium
Ag^+ Silver Cation
OH^- Hydroxide
NO3^- Nitrate
HCO3^- Hydrogen Carbonate
SO4^2- Sulfate
CO3^2- Carbonate
PO4^3- Phosphate
what is a polyatomic ion
a polyatomic ion is an ion composed of more than one atom
factors affecting bond strength (bond enthalpies)
1) charge: In ionic compounds, the greater the charge of the ions, the stronger the electrostatic attraction between them, leading to higher bond enthalpy
2) bond order: Higher bond order (single, double, triple bonds) increases bond enthalpy because more electrons are shared between the atoms. this results in stronger and shorter bonds
3) atomic radius: Smaller atoms tend to form stronger bonds due to shorter bond lengths, which increase the attractive force between the nuclei and bonding electrons
Diatomic elements
H2
C
Si
N2
O2
F2
Cl2
Br2
I2
P4
S8
endothermic
An endothermic reaction is a chemical reaction that absorbs energy from its surroundings, usually in the form of heat. This results in a decrease in the temperature of the surroundings and is characterised by a positive change in enthalpy (ΔH>0)
exothermic
An exothermic reaction is a chemical reaction that releases energy, usually in the form of heat, to its surroundings. This results in an increase in temperature of the surroundings and can be represented by a negative change in enthalpy (ΔH<0)
Determining bonding type: metallic character and effective nuclear charge (Zeff)
the weaker an atom attracts electrons the more metallic it is. the stronger an atom attracts electrons, the more nonmetallic it is.
Determining bonding type: electronegativity difference
greater electronegativity difference between the bonding atoms leads to greater and greater polarity in terms of how the electrons are attracted/spread out between atoms.
in covalent bonding, a smaller electronegativity difference leads to the sharing of electrons between atoms. With very small differences this sharing very even and referred to as nonpolar. as the difference increases, the electron sharing becomes more uneven and increasingly polar.
in ionic bonds, the electronegativity difference is so great that electrons are completely transferred from one atom to another, resulting in the formation of charged ions (positive and negative ions)
in metallic bonds the EN values are so low (they attract electrons so weakly) that they result in electron delocalisation.
polarity
how evenly electron pairs are shared. polarity depends on the electronegativity difference of the bonding atoms.
electronegativity (χ) difference formula
Δχ = |χa - χb|
average electronegativity formula
∑χ = (χa + χb) / 2
bond enthalpies (kJ/mol)
the energy required to break (1 mole of) a bond in the gas phase.
for covalent compounds
bond breaking involves the separation of particles that are attracted to each other, therefore it is always an endothermic process. bond formation is the opposite process and therefore results in an equal magnitude, but opposite enthalpy change.
bond enthalpy can also be seen as a measure of bond strength.
Explain the melting point trend down group 1.
n Group 1 (alkali metals), the melting points decrease as you move down the group, from lithium (Li) to cesium (Cs).
due to:
Increasing atomic/ionic radii of metal
Decreasing electrostatic attraction to delocalized electrons/ decreasing metallic bond strength
Explain the melting point trend from sodium, to magnesium, to aluminum.
In Period 3, the melting points increase as you move from sodium (Na) to magnesium (Mg) to aluminum (Al).
due to:
Increasing ionic charge on metal cation AND decreasing ionic/atomic radii of metal ion
Increasing electrostatic attraction to delocalized electrons/ increasing metallic bond strength
Explain the melting points of silicon and phosphorus compared to those of the period 3 metals.
Silicon is a giant covalent (network) solid. Only covalent bonds between atoms. Extremely strong bonding and very high melting point compared to most metals.
Phosphorus is molecular. Weak intermolecular forces between molecules. Very low melting point compared to most metals.
lattice enthalpy (kJ/mol)
Lattice enthalpy is the amount of energy required to separate one mole of an ionic solid into its gaseous ions or, conversely, the energy released when gaseous ions come together to form one mole of an ionic solid.
for ionic compounds
A more negative lattice enthalpy indicates a more stable ionic compound because strong ionic bonds result in more energy being released when the solid forms from gaseous ions.
It reflects the strength of the ionic interactions in the crystal lattice.
