Introductory Chemistry – Atoms and Chemical Bonds (Lecture 01)

Question-and-answer flashcards covering atomic structure, electron configuration, metallic, ionic and covalent bonding, hydrogen bonding, polarity, solubility rules, polyatomic ions, hydrates, and related properties.

What is the simplest identifiable particle of an element?

An atom.

State two characteristics that are identical for atoms of the same element.

They have the same size and mass.

Approximately how wide is an atom?

About 0.1 nanometre.

What does the atomic (proton) number represent?

The number of protons in the nucleus (and, in a neutral atom, the number of electrons).

How is nucleon number calculated?

Nucleon number = number of protons + number of neutrons.

Give the maximum number of electrons that can occupy the first, second and third shells (for the first 20 elements).

1st shell: 2; 2nd shell: 8; 3rd shell: 8.

What is meant by ‘valence electrons’?

Electrons in the outermost electron shell of an atom.

How are elements arranged in the periodic table?

In increasing order of proton (atomic) number.

What do elements in the same period share?

They have the same number of electron shells.

What do elements in the same group share?

They have the same number of valence electrons.

How does a positive ion form?

By an atom losing one or more electrons.

How does a negative ion form?

By an atom gaining one or more electrons.

Define metallic bonding.

Attraction between positive metal ions and a ‘sea’ of delocalised electrons.

List two physical properties of metals that arise from metallic bonding.

High melting/boiling points and electrical conductivity in solid and molten states.

Define ionic bonding.

The electrostatic attraction between oppositely charged ions, usually formed from metal and non-metal atoms via electron transfer.

Why do ionic compounds conduct electricity when molten or in solution but not when solid?

Because ions are free to move in molten/aqueous states but fixed in solid lattices.

Give two general properties of ionic compounds.

High melting/boiling points and (many) solubility in water but not in most organic solvents.

Define covalent bonding.

Attraction between two nuclei and a shared pair of electrons (electron sharing between non-metals).

Why do small covalent molecules have low melting and boiling points?

Only weak intermolecular (van der Waals’) forces need to be overcome, despite strong intramolecular covalent bonds.

Do small covalent molecules conduct electricity? Why/why not?

No, because they lack free ions or delocalised electrons.

What is a hydrogen bond?

An intermolecular attraction between a hydrogen atom bonded to a highly electronegative atom (e.g., O, N, F) and a lone pair on another electronegative atom.

Define electronegativity.

The ability of an atom in a molecule to attract shared electrons toward itself.

What is a polar covalent bond?

A covalent bond with unequal electron sharing due to a difference in electronegativity, creating partial charges (δ+ and δ−).

State the condition for a molecule to be polar.

It must have polar bonds and a molecular geometry that does not cancel the dipole moments, giving a permanent dipole.

Why is water able to dissolve many ionic compounds?

Water’s polarity allows it to surround and separate positive and negative ions, overcoming the ionic lattice forces.

Explain why 2-butanol is more soluble in water than hexane.

2-Butanol has an –OH group that forms hydrogen bonds with water (polar), whereas hexane is non-polar and lacks such interactions.

State the ‘like dissolves like’ rule for solubility.

Polar (or ionic) solutes dissolve best in polar solvents; non-polar solutes dissolve best in non-polar solvents.

What are polyatomic ions? Give two examples.

Ions composed of two or more atoms covalently bonded with an overall charge; e.g., NH4⁺ (ammonium), SO4²⁻ (sulfate).

What is water of crystallisation?

Water molecules incorporated into the crystal lattice of some ionic compounds (hydrates).

How would you name CuSO4·5H2O using hydrate nomenclature?

Copper(II) sulfate pentahydrate.

Give the hydrate prefix for six water molecules.

Hexa- (e.g., hexahydrate).

Why is calcium carbonate insoluble in water whereas sodium carbonate is soluble?

The ionic attraction between Ca²⁺ and CO3²⁻ is stronger than between Na⁺ and CO3²⁻; all Group 1 metal salts (like Na⁺) are water-soluble.

Is CO2 polar or non-polar? Explain.

Non-polar; although C=O bonds are polar, the linear molecule’s symmetric shape cancels dipoles.

What types of substances typically dissolve in organic solvents like hexane or benzene?

Non-polar substances (e.g., hydrocarbons, fat-soluble vitamins).

What causes metals to be malleable and ductile?

The ability of metal ions to slide over one another while remaining bonded by the sea of delocalised electrons.

Describe the difference between intramolecular and intermolecular forces.

Intramolecular forces are within molecules (e.g., covalent bonds); intermolecular forces occur between molecules (e.g., van der Waals’, hydrogen bonds).