Topic 2: Atomic Structure and the Periodic Table

Protons identify the element. Neutrons add nuclear mass without changing the element's identity. Electrons are attracted to the positive nucleus and are arranged in energy levels and sub-shells; their arrangement later explains bonding and chemical behaviour.

A neutral atom has equal numbers of protons and electrons, so its positive and negative charges cancel. Changing the number of electrons makes an ion; changing the number of neutrons makes another isotope of the same element.

An electric field deflects protons and electrons because they are charged, but it does not deflect neutrons. Electron mass is so small compared with nucleon mass that mass number counts protons and neutrons only.

In $^{84}_{36}\ce{Kr}$, $Z=36$, so the element is krypton and its nucleus contains 36 protons. $A=84$ means the nucleus contains 84 nucleons in total. Mass number is always a whole number for one isotope.

Here $q$ is the ion charge in elementary-charge units: $q=+2$ for a $2+$ ion and $q=-1$ for a $1-$ ion. Subtracting a negative charge therefore adds an electron.

$^{35}\ce{Cl}$ and $^{37}\ce{Cl}$ both contain 17 protons. They contain 18 and 20 neutrons respectively. Neutral atoms of both isotopes contain 17 electrons and therefore have the same ground-state electronic configuration.

Because chemical reactions mainly involve electrons, isotopes of an element have very similar chemical properties. Their different masses can produce different physical behaviour and separate peaks in a mass spectrum.

1. Introduce and vaporise the sample where necessary.
2. Ionise particles to form positive ions.
3. Accelerate ions using an electric field.
4. Separate ions because different $m/z$ values respond differently or take different flight times.
5. Detect ions and convert the signal into a spectrum.

Elemental peaks reveal isotope masses and abundances; the weighted mean gives $A_r$, and the equation can be rearranged to find an unknown abundance or mass. For a molecule, the molecular-ion peak gives its relative molecular mass when $z=1$, helping identify the molecule. Fragment peaks represent smaller ions.

For two different isotopes, count both arrangements: $^{35}\ce{Cl}$-$^{37}\ce{Cl}$ and $^{37}\ce{Cl}$-$^{35}\ce{Cl}$. This factor of two is why the mixed-isotope peak is larger than either single arrangement alone.

The equations define first, second and third ionisation energies respectively. Values are quoted per mole, usually in $\mathrm{kJ\,mol^{-1}}$, and the starting species must be gaseous in every case.

An empty orbital holds no electrons; a singly occupied orbital holds one; a full orbital holds a pair. When two electrons share an orbital, their spins must be opposite, represented in electron-in-box notation by arrows pointing in opposite directions, $\uparrow\downarrow$.

Orbitals are grouped into sub-shells. An s sub-shell contains one orbital, a p sub-shell three orbitals and a d sub-shell five orbitals. Orbital capacity therefore controls sub-shell capacity.

A 2p electron is higher in energy and more shielded than a 2s electron, explaining why boron's first ionisation energy is lower than beryllium's despite boron having more protons. Within 2p, repulsion in a paired orbital makes an electron easier to remove from oxygen than from nitrogen's singly occupied 2p orbitals.

For one element, successive ionisation energies rise, but a very large jump occurs after all electrons in the outer shell have been removed. The next electron comes from an inner shell, closer to the nucleus and less shielded. For a main-group element, the number of electrons removed before the first large jump identifies its group.

Across successive elements, first ionisation energy generally rises within a period as nuclear charge increases. A large fall after a noble gas shows the start of a new shell. Smaller falls, such as Be to B and Mg to Al, show entry into a higher-energy p sub-shell; the N to O and P to S falls show the effect of pairing within p orbitals.

Each p sub-shell contains three orbitals of the same basic shape and energy, oriented along mutually perpendicular axes. They are labelled $p_x$, $p_y$ and $p_z$. Each orientation is a different orbital and can hold up to two opposite-spin electrons.

The boundary shape represents a region of high probability for locating an electron; it is not a solid surface. The two p lobes belong to one orbital, not two separate orbitals.

Single occupation keeps electrons apart while the orbitals have equal energy. When two electrons share an orbital, opposite spins are required. The order of equivalent boxes does not matter; the occupancy pattern does.

Use superscripts for electron counts: s holds 2, p holds 6 and d holds 10. For example, sulfur is $1s^2 2s^2 2p^6 3s^2 3p^4$; electron-in-box notation shows one paired and two singly occupied 3p orbitals. Noble-gas shorthand may replace completed inner shells.

Chromium and copper use the observed configurations $[\ce{Ar}]3d^5 4s^1$ and $[\ce{Ar}]3d^{10}4s^1$. For ions, add electrons for negative charge and remove them for positive charge from the highest principal shell first: transition-metal ions lose 4s electrons before 3d electrons. Thus $\ce{Fe^{3+}}$ is $[\ce{Ar}]3d^5$ and $\ce{Cl-}$ is $[\ce{Ar}]$.

Elements in the same group have the same pattern of outer-shell electrons. Group 1 atoms have an outer $ns^1$ electron and commonly form $1+$ ions by losing it. Group 17 atoms have outer $ns^2np^5$ configurations and need one more electron for a filled outer shell, so they commonly form $1-$ ions or one covalent bond.

Moving across a period changes the outer configuration one electron at a time, producing a systematic change from metallic to non-metallic behaviour. Moving down a group preserves the valence pattern but adds shells, so chemical behaviour remains related while reactivity can change.

In the first four quantum shells, the relevant sub-shells are $1s$; $2s,2p$; $3s,3p,3d$; and $4s,4p,4d$. Each s, p or d sub-shell always has the same capacity regardless of its shell number.

Use atomic number on the horizontal axis so the elements are in sequence. Label the vertical axis with the property and its unit, choose a scale that uses the plotting area effectively, plot each value accurately, and join points only when the purpose is to display the sequence rather than imply continuous values between elements.

For a logarithmic treatment, calculate $\log_{10}$ of each first-ionisation-energy value and plot that transformed value against atomic number. Label the axis to show both the logarithm and the original energy unit convention. Equal vertical intervals then represent equal ratios in first ionisation energy, not equal raw-energy differences, which can make repeated proportional features easier to compare.

Across each period, first ionisation energy generally increases because proton number rises while added electrons enter the same main shell, so shielding changes little and atomic radius decreases. The Group 13 dip occurs when the electron removed is in a higher-energy p sub-shell; the Group 16 dip occurs because repulsion in a paired p orbital makes removal easier than from the preceding half-filled p sub-shell.

Down a group, first ionisation energy decreases. The outer electron occupies a higher shell, farther from the nucleus, and experiences more inner-shell shielding. These effects outweigh the increased proton number, weakening attraction to the electron removed.