Chemical Bonding & Structures – Lecture Review

A comprehensive set of Q&A flashcards covering definitions, bonding types, structures, properties, VSEPR, and examples for exam revision.

What is the definition of an element?

A pure substance that cannot be broken down into simpler substances by chemical methods.

How is a compound defined?

A pure substance containing two or more elements chemically combined in a fixed ratio.

What is a mixture?

A combination of elements and/or compounds that are not chemically combined and have no fixed ratio.

How is a compound formed compared to a mixture?

Compounds form through chemical reactions; mixtures form through physical mixing.

Do mixtures have fixed melting and boiling points?

No. They melt and boil over a range of temperatures.

What is the octet rule?

Atoms achieve stability when they have eight electrons in their valence shell (two if it is the first shell).

How can you quickly identify a metal versus a non-metal based on valence electrons?

1–3 valence electrons → metal; 5–7 valence electrons → non-metal; 4 can be either.

Define an ion.

A charged particle formed when an atom gains or loses electrons.

What charge does a cation carry and how is it formed?

Positive; formed when an atom (usually a metal) loses electrons.

What charge does an anion carry and how is it formed?

Negative; formed when an atom (usually a non-metal) gains electrons.

Describe an ionic bond.

The strong electrostatic force of attraction between oppositely charged ions (metal + non-metal).

What structure do ionic compounds form?

A giant ionic crystal lattice of alternating positive and negative ions.

Why do ionic substances have high melting and boiling points?

Large amounts of energy are required to overcome strong electrostatic forces between ions in the lattice.

Why are ionic compounds brittle?

When layers shift under stress, like-charged ions align, repel, and the lattice shatters.

In which states do ionic compounds conduct electricity and why?

Molten or aqueous; ions are mobile and can carry charge. They do not conduct when solid.

Are most ionic substances soluble in water or organic solvents?

Generally soluble in water and insoluble in organic solvents.

Define a covalent bond.

A bond formed when two non-metal atoms share one or more pairs of electrons.

What is the difference between a single, double, and triple covalent bond?

Single shares 1 pair; double shares 2 pairs; triple shares 3 pairs of electrons.

What are simple molecular substances?

Discrete molecules held together by weak intermolecular forces; most covalent substances fall here.

Why do simple molecular substances have low melting and boiling points?

Only weak intermolecular forces need to be overcome, requiring little energy.

Do simple molecular substances conduct electricity?

No, because they lack free-moving electrons or ions.

Define a giant covalent (network) structure.

A vast lattice where atoms are linked by extensive covalent bonds (e.g., diamond, SiO2, graphite).

Why is diamond hard while graphite is soft?

Diamond has a 3-D tetrahedral network of strong covalent bonds; graphite has layers held by weak forces that slide easily.

Why does graphite conduct electricity but diamond does not?

Graphite has delocalised electrons between layers; diamond has no free electrons.

What is a metallic bond?

Strong electrostatic attraction between positive metal ions and a ‘sea’ of delocalised electrons.

Why are metals malleable and ductile?

Regular layers of ions can slide over each other when force is applied.

Why do metals have high melting and boiling points?

Strong metallic bonds require large amounts of energy to break.

Why are alloys harder than pure metals?

Different-sized atoms disrupt the regular layers, preventing them from sliding easily.

What is a polyatomic ion?

An ion composed of two or more covalently bonded atoms acting as a single charged species (e.g., SO₄²⁻).

Give an example of a common cation with a +2 charge.

Mg²⁺ (magnesium ion).

Give an example of a common anion with a –1 charge.

Cl⁻ (chloride ion).

State three physical properties used to distinguish bonding/structures.

Melting/boiling point, electrical conductivity, solubility.

Which structures conduct electricity in all states?

Metals and graphite.

Which type of substances typically dissolve in organic solvents?

Simple molecular (covalent) substances and many metals/alloys.

What does VSEPR theory state?

Electron pairs around a central atom repel to positions as far apart as possible, determining molecular shape.

Why is the H–O–H bond angle in water 104.5° instead of 109.5°?

Two lone pairs on oxygen repel more strongly than bond pairs, compressing the bond angle.

Explain expanded octet and give one example.

Atoms from Period 3 onward can hold more than eight valence electrons, e.g., SF₆ where sulfur has 12 electrons.

What is coordinate (dative) bonding?

A covalent bond where both shared electrons originate from the same atom.

Why do higher ionic charges increase melting points?

Greater charge produces stronger electrostatic attraction, requiring more energy to break.

Describe the conductivity difference between ionic compounds and metals.

Ionic compounds conduct via mobile ions in liquid/aqueous states; metals conduct via delocalised electrons in all states.

What property identifies a substance as a mixture rather than a pure substance?

It can be separated by physical methods and has variable composition.

Why do polymers soften over a range of temperatures rather than melt sharply?

They have varying chain lengths and weaker intermolecular forces that overcome gradually.

What determines whether a covalent molecule is simple or giant?

The number of atoms and whether the covalent bonding extends throughout a continuous network.

Describe the bonding and structure in silicon dioxide (SiO₂).

Each Si atom covalently bonds to four O atoms in a tetrahedral network, forming a giant covalent lattice.

Explain why metals are good conductors of heat.

Delocalised electrons transfer kinetic energy quickly throughout the lattice.

Which charge carriers are present in molten NaCl during conduction?

Mobile Na⁺ and Cl⁻ ions.

What is malleability?

The ability of a substance to be hammered into thin sheets without breaking.

How does the presence of lone pairs affect molecular geometry?

Lone pairs exert greater repulsion than bond pairs, altering bond angles and overall shape.

Which bonding type involves ‘loss of electrons’ by atoms to form a lattice with delocalised electrons?

Metallic bonding.

Name two alloys and their constituent elements.

Brass (copper + zinc); Bronze (copper + tin).