AQA A Level Chemistry - Molecules: Shapes & Forces

Lone pairs = Repel each other more than bonded pairs Bond pairs = Represent shared electrons between atoms Double bonds = Treated the same as single bonds in repulsion Triple bonds = Have greater repulsion than double bonds

Valence shell electrons = Found in the outer shell of an atom Electron repulsion = Same charge repels other electrons Lone pair effect = Repels more than bonded pairs Stable shape = Adopted to minimize repulsion forces

Linear = 2 bonding pairs, 0 lone pairs Trigonal planar = 3 bonding pairs, 0 lone pairs Bent = 2 bonding pairs, 1 lone pair Tetrahedral = 4 bonding pairs, 0 lone pairs

<p>Lone pair - Lone pair = Strongest repulsion Lone pair - Bond pair = Intermediate repulsion Bond pair - Bond pair = Weakest repulsion Double bond - Double bond = Greater than single bonds</p> Signup and view all the answers

<p>Valence shell = Outermost shell of an atom Repulsion = Forces between like charges Concentrated charge cloud = Typically associated with lone pairs Molecule stability = Minimized repulsive forces between pairs</p> Signup and view all the answers

<p>Bonded pairs = Electrons shared between atoms Lone pairs = Electrons localized on a single atom Repulsive forces = Influence the molecular shape Electron pairs = Determine bond angles in molecules</p> Signup and view all the answers

<p>Single bond = One pair of shared electrons Double bond = Two pairs of shared electrons Triple bond = Three pairs of shared electrons Multiple bonds = Involve increased electron repulsion</p> Signup and view all the answers

<p>Electron pair arrangement = Affects molecular geometry Bond angles = Determined by electron pair repulsion Shape stability = Minimized repulsion leads to this VSEPR theory = Predicts shapes based on electron pairs</p> Signup and view all the answers

<p>Lone pairs = Higher repulsive force Bonding pairs = Lower repulsive force Double bonds = Intermediate repulsive force Triple bonds = Higher repulsive force than single bonds</p> Signup and view all the answers

<p>CCl4 = Tetrahedral N(CH3)3 = Trigonal pyramidal Phosphorus(V) chloride = Trigonal bipyramidal CO2 = Linear</p> Signup and view all the answers

<p>Tetrahedral = $109.5^{ ext{o}}$ Trigonal planar = $120^{ ext{o}}$ Trigonal bipyramidal = $90^{ ext{o}}$ and $120^{ ext{o}}$ Linear = $180^{ ext{o}}$</p> Signup and view all the answers

<p>Group 15 = 5 valence electrons Group 17 = 7 valence electrons Group 14 = 4 valence electrons Group 16 = 6 valence electrons</p> Signup and view all the answers

<p>Trigonal bipyramidal = Phosphorus(V) chloride Tetrahedral = CCl4 Trigonal pyramidal = N(CH3)3 Linear = CO2</p> Signup and view all the answers

<p>Lone pair = Non-bonding electron pair Bonding pair = Electron pair shared between atoms Electronegativity = Ability to attract electrons in a bond Molecular geometry = 3D arrangement of atoms in a molecule</p> Signup and view all the answers

<p>VSEPR = Theory to predict molecular shapes Covalent bond = Sharing of electron pairs between atoms Valence electrons = Electrons in the outermost shell Bond angle = Angle between two bonds in a molecule</p> Signup and view all the answers

<p>Single bond = C–H in methane Double bond = O=O in oxygen Triple bond = N≡N in nitrogen Polar covalent bond = H–Cl in hydrochloric acid</p> Signup and view all the answers

<p>Sodium = Higher electronegativity than caesium Caesium = Lower electronegativity than sodium Oxygen = Higher electronegativity than nitrogen Bromine = Lower electronegativity than chlorine</p> Signup and view all the answers

<p>Groups = Electronegativity decreases down the group Periods = Electronegativity increases across the period Transition Metals = Variable electronegativity Noble Gases = Generally low electronegativity</p> Signup and view all the answers

<p>Increasing nuclear charge = Increases electronegativity Adding new electron shells = Decreases electronegativity Increasing atomic radii = Decreases electronegativity Constant shielding across a period = Increases electronegativity</p> Signup and view all the answers

<p>High electronegativity = Strong attraction for bonding electrons Low electronegativity = Weak attraction for bonding electrons Increased shielding = Less attraction between nucleus and outer electrons Decreased shielding = Greater attraction between nucleus and outer electrons</p> Signup and view all the answers

<p>Electronegativity trend in groups = Decreases due to increased shielding Electronegativity trend across periods = Increases due to higher nuclear charge Large atomic radii down a group = Due to more filled electron shells Small atomic radii across a period = Due to increased attraction between nucleus and electrons</p> Signup and view all the answers

<p>Outer electrons = Experience attraction from the nucleus Inner electrons = Shield the outer electrons Valence electrons = Participate in chemical bonding Core electrons = Do not participate in bonding</p> Signup and view all the answers

<p>Sodium = Group 1 Chlorine = Group 17 Bromine = Group 17 Calcium = Group 2</p> Signup and view all the answers

<p>Intramolecular forces = Forces within a molecule, usually covalent bonds Covalent bonds = Bonds formed by sharing outer electrons between atoms Intermolecular forces = Forces between molecules that are weaker than intramolecular forces Hydrogen bonding = A special type of permanent dipole - dipole force</p> Signup and view all the answers

<p>Induced dipole - dipole forces = Also called van der Waals or London dispersion forces Permanent dipole - dipole forces = Attractive forces between neighboring molecules with a permanent dipole Hydrogen bonding = Stronger than standard dipole - dipole interactions Intermolecular forces = Weaker forces compared to intramolecular forces</p> Signup and view all the answers

<p>Fluorine = Most electronegative element in the periodic table Electronegativity = Increases going across periods in the periodic table Top right of the periodic table = Region with highest electronegativity Electronegativity trend = General increase towards the top right corner</p> Signup and view all the answers

<p>Covalent bond = Stronger than hydrogen bonds Hydrogen bond = About one tenth the strength of covalent bonds Intramolecular forces = Stronger than intermolecular forces Intermolecular forces = Weaker than intramolecular forces</p> Signup and view all the answers

<p>Single bond = Involves one shared pair of electrons Double bond = Involves two shared pairs of electrons Triple bond = Involves three shared pairs of electrons Coordinate bond = Involves one atom donating both electrons</p> Signup and view all the answers

<p>Intramolecular forces = Covalent bonds within a molecule Intermolecular forces = Attractive forces between different molecules Electronegativity = Tendency of an atom to attract electrons Lewis structures = Diagrams showing covalent bonds and molecule structure</p> Signup and view all the answers

<p>Induced dipole forces = Temporary shifts in electron distribution in molecules Permanent dipole-dipole forces = Ongoing attractions between permanent dipoles Hydrogen bonding = Special strong interaction involving hydrogen atoms Dipole-induced dipole forces = One polar molecule inducing a dipole in another</p> Signup and view all the answers

<p>Increasing electronegativity = Represents increasing attraction for electrons Periodic table trends = Electronegativity increases across periods Fluorine's position = Indicates it has the highest electronegativity Atoms with higher electronegativity = Tend to attract electrons more strongly</p> Signup and view all the answers

<p>Nonpolar bond = Electrons shared equally Polar bond = Electrons drawn toward more electronegative atom Permanent dipole = Always has a negative and positive charge end Asymmetric electron distribution = Charge centers do not coincide</p> Signup and view all the answers

<p>Electronegativity = Ability of an atom to attract electrons δ+ (delta positive) = Partial charge on less electronegative atom δ- (delta negative) = Partial charge on more electronegative atom Dipole-dipole forces = Attraction between polar molecules due to permanent dipoles</p> Signup and view all the answers

<p>Chlorine (Cl) = Higher electronegativity than hydrogen (H) Hydrogen (H) = Lower electronegativity than chlorine (Cl) Oxygen (O) = Significantly high electronegativity Carbon (C) = Moderate electronegativity compared to metals</p> Signup and view all the answers

<p>Greater difference in electronegativity = More polar the bond Same electronegativities = Forms a nonpolar bond Permanent dipole = Induces attraction between molecules Asymmetry in electron distribution = Differentiates between polar and nonpolar bonds</p> Signup and view all the answers

<p>Polar bond = HCl Nonpolar bond = Cl2 Permanent dipole = Water (H2O) Covalent bond = O2</p> Signup and view all the answers

<p>Nonpolar covalent bond = Equal sharing of electrons Polar covalent bond = Unequal sharing of electrons Ionic bond = Transfer of electrons Metallic bond = Sea of delocalized electrons</p> Signup and view all the answers

<p>Permanent dipole = Indicates polarity of molecules Asymmetrical charge distribution = Leads to dipole interactions Polar molecule = Has a dipole moment Nonpolar molecule = Does not produce a net dipole moment</p> Signup and view all the answers

<p>Cl-H bond = Cl attracts electrons more H-H bond = Electrons shared equally O-H bond = O draws electrons significantly C-C bond = Equal attraction in sigma bond</p> Signup and view all the answers

<p>Oxygen = Highly electronegative atom that can bond with hydrogen Hydrogen = Becomes δ+ charged when bonded to O, N or F Nitrogen = Another electronegative atom involved in hydrogen bonding Fluorine = The most electronegative element capable of hydrogen bonding</p> Signup and view all the answers

<p>High boiling point = Strong hydrogen bonding High surface tension = Intermolecular forces Higher density in liquid than solid = Hydrogen bond structure in ice Anomalous properties = Presence of hydrogen bonds</p> Signup and view all the answers

<p>Hydrogen bonds = Strong bonds between molecules with hydrogen and O, N or F Permanent dipole – permanent dipole = Interaction between polar molecules Dispersion forces = Induced dipole interactions due to electron movement Covalent bonds = Strong bonds within a molecule</p> Signup and view all the answers

<p>Solid H2O = Hydrogen bonds maintain a fixed structure Liquid H2O = Molecules can move but are still held together by hydrogen bonds Ice = Lower density due to expanded hydrogen-bonded structure Steam = Higher energy state where hydrogen bonds are broken</p> Signup and view all the answers

<p>Enthalpy of vaporization = Energy required to boil water Melting point = Energy needed to transition from solid to liquid Boiling point = Temperature at which water vaporizes Surface tension = Resistance of water to increase surface area</p> Signup and view all the answers

<p>O = Forms two hydrogen bonds N = Bonds with δ+ hydrogen F = Highly reactive due to electronegativity H = Participates in hydrogen bonding but not as an acceptor</p> Signup and view all the answers

<p>High melting point = Comparison to other simple molecules High boiling point = Increased energy requirement due to hydrogen bonds High surface tension = Significant impact on water's interaction with other substances Anomalous properties = Unique behavior due to hydrogen bonding</p> Signup and view all the answers

<p>Strong intermolecular forces = Explanation for unusual properties of water Bonding angles = Influence on molecular geometry Lone pairs = Necessary for forming hydrogen bonds Polarization of hydrogen = Allows for bond formation with electronegative atoms</p> Signup and view all the answers

<p>Water (H2O) = Can form two hydrogen bonds Ammonia (NH3) = Can form one hydrogen bond per molecule Hydrogen fluoride (HF) = Capable of very strong hydrogen bonds Methane (CH4) = Cannot form hydrogen bonds at all</p> Signup and view all the answers

Valence Shell Electron Pair Repulsion theory predicts the shapes and bond angles of molecules.

Negatively charged electrons in a molecule that repel each other, influencing molecular shape.

Electrons in the outermost shell of an atom, influencing bonding and molecular shape.

Repulsion between lone electron pairs is stronger than between bonding pairs, affecting molecular shape.

Repulsion between bonded electron pairs in a molecule.

The three-dimensional arrangement of atoms in a molecule influenced by the repulsive forces of electron pairs.

The angle between two bonded atoms around a central atom, affected by electron repulsion.

Interaction of negative charges (electrons) leading to the minimization of repulsive forces, influencing molecular shapes and bond angles.

The angle between two bonds emanating from the same atom.

A molecular shape with five atoms or groups of atoms bonded to a central atom, arranged in a specific three-dimensional geometry.

A molecule with phosphorus as the central atom and five chlorine atoms bonded to it.

Electron pairs around a central atom repel each other, leading to specific molecular shapes.

Chemical bonds formed by sharing electrons between atoms.

Inner electrons prevent outer electrons from fully experiencing the full positive charge of the nucleus.

Electronegativity decreases as you move down a group on the periodic table.

Electronegativity increases as you move across a period on the periodic table.

Increased shielding decreases electronegativity; decreased shielding increases electronegativity.

Increased shielding leads to a larger atomic radius.

The positive charge of the atom's nucleus, determined by the number of protons.

The tendency of an atom to attract bonding electrons towards itself.

Half the distance between the nuclei of two bonded atoms.

Electronegativity increases as you move across a period on the periodic table.

Fluorine is the most electronegative element.

A force within a molecule, typically a covalent bond.

A force between molecules. Weaker than intramolecular forces.

A type of intermolecular force (van der Waals/London dispersion forces).

Attractive forces between polar molecules with permanent dipoles.

A very strong type of permanent dipole-dipole force.

A covalent bond where electrons are unequally shared between atoms due to differences in electronegativity.

A covalent bond where electrons are shared equally between atoms with the same electronegativity.

A slight positive (δ+) or negative (δ-) charge on an atom in a polar bond.

The covalent bond becomes polar, forming partial positive and negative charges on the atoms. The less electronegative atom gets a δ+, and the more electronegative atom gets a δ-.

A molecule with a permanent separation of positive and negative charges due to polar bonds.

Attractive forces between polar molecules due to the interaction of their permanent dipoles.

The greater the difference in electronegativity, the more polar the bond becomes.

Attractive forces that exist between molecules, holding them together.

Hydrogen bonding occurs when hydrogen is bonded to a highly electronegative atom (O, N, or F). This creates a highly polarized bond, making the hydrogen atom very δ+ charged. It can then form a strong interaction with a lone pair on another electronegative atom.

Due to strong hydrogen bonding, water exhibits unusual properties compared to other similar molecules, such as exceptionally high melting and boiling points, high surface tension, and a higher density in liquid form compared to solid form.

Water's high melting and boiling points are a result of the strong hydrogen bonds between its molecules. Breaking these bonds requires significant energy.

The amount of energy required to convert a liquid into a gas at a constant pressure.

Weak intermolecular forces caused by temporary, induced dipoles in nonpolar molecules due to fluctuations in electron distribution.

Comparing enthalpy changes of vaporization in a series of similar molecules helps to understand the relative strength of intermolecular forces. For example, comparing H2S, H2Se, and H2Te shows increasing enthalpy changes as the molecules get larger due to stronger dispersion forces. However, H2O deviates significantly due to its strong hydrogen bonding.

Hydrogen bonding is responsible for many of water's essential properties, including its role as a solvent, its ability to moderate temperature, and its properties as a transport medium.

When drawing hydrogen bonds, use a dashed, straight line connecting the lone pair of the electronegative atom to the δ+ hydrogen atom. Make it clear where the bond starts and ends.