Atomic Structure and Properties

Because atoms are far too small to count, chemists count in moles 摩尔. One mole is Avogadro's number 阿伏伽德罗常数 of particles, $N_A=6.022\times10^{23}$. The molar mass 摩尔质量 (grams per mole, read off the periodic table) converts between mass and moles:

Moles are the bridge between the lab (grams you weigh) and the equation (particles that react).

Worked example. How many moles, and how many molecules, are in $36.0\ \text{g}$ of water ($M=18.0\ \text{g/mol}$)?

A mass spectrometer 质谱仪 separates atoms by mass, giving a mass spectrum 质谱: peaks at each isotope 同位素 (same element, different neutron count) with heights showing their relative abundance 相对丰度. The element's average atomic mass is the abundance-weighted average of its isotope masses.

Worked example. Chlorine is $75.8\%$ chlorine-35 and $24.2\%$ chlorine-37. Its average atomic mass is

The average lies closer to $35$ because that isotope is more abundant – which is why the periodic-table value is $35.5$, not a whole number.

The percent composition 百分组成 of a compound is each element's mass fraction. From it you find the empirical formula 实验式 (simplest whole-number ratio of atoms) by converting each element's mass to moles and dividing by the smallest. The molecular formula 分子式 is a whole-number multiple of the empirical formula, found from the molar mass.

Worked example. A compound is $40.0\%$ C, $6.7\%$ H, and $53.3\%$ O by mass. Assuming $100\ \text{g}$, convert each mass to moles and divide by the smallest:

so the empirical formula is $\text{CH}_2\text{O}$. If the molar mass were $180\ \text{g/mol}$ ($=6\times30$), the molecular formula would be $\text{C}_6\text{H}_{12}\text{O}_6$ – glucose.

Unlike a pure compound, a mixture 混合物 has variable composition – its parts keep their own identities. Describe a mixture by the mass or mole fraction of each component; these do not follow a fixed formula. Spectroscopy (like PES or absorption) can measure how much of each component is present.

An atom is a tiny nucleus 原子核 (protons and neutrons) surrounded by electrons in shells and subshells (s, p, d, f). The electron configuration 电子排布 lists how electrons fill these, lowest energy first (e.g. $1s^2\,2s^2\,2p^6$). The outermost, highest-energy electrons – the valence electrons 价电子 – control chemistry. Coulomb's law explains their energies: electrons closer to, and less shielded from, the nucleus are held more tightly.

Worked example. Write the electron configuration of sulfur ($Z=16$). Fill subshells in order until $16$ electrons are placed: $1s^2\,2s^2\,2p^6\,3s^2\,3p^4$. The $3s$ and $3p$ electrons (six in total) are the valence electrons, so sulfur tends to gain two electrons to complete its octet.

Photoelectron spectroscopy 光电子能谱 (PES) measures the energy needed to remove electrons from each subshell. Each peak is a subshell: its position gives the binding energy (how tightly held) and its height gives the number of electrons in it. PES data let you read an element's electron configuration and confirm shell structure directly.

Trends across the periodic table follow from nuclear charge and shielding:

• Atomic radius 原子半径 decreases across a period (stronger pull) and increases down a group (more shells).
• Ionization energy 电离能 (energy to remove an electron) increases across, decreases down – opposite to radius.
• Electronegativity 电负性 (pull on shared electrons) increases across and up, toward fluorine.

Atoms gain, lose, or share valence electrons to reach stable configurations. Metals (low ionization energy) lose electrons to form cations 阳离子; nonmetals gain electrons to form anions 阴离子. Oppositely charged ions attract into an ionic compound 离子化合物, whose formula balances the charges to make the whole neutral. For example, aluminium ($3+$) and oxide ($2-$) combine as $\text{Al}_2\text{O}_3$ so the $+6$ and $-6$ cancel.

• Use the mole as the hub: convert grams↔moles with $n=m/M$ and moles↔particles with Avogadro's number.
• Relative atomic mass is the abundance-weighted average of the isotopes — it lies closer to the more abundant one, not halfway.
• For an empirical formula, turn each element's mass into moles and divide by the smallest; scale up to whole numbers.
• Read periodic trends from nuclear charge and shielding: atomic radius decreases across a period, ionisation energy and electronegativity increase across (and up).
• Fill electron configurations in energy order; the outer valence electrons control the chemistry.