How is the periodic table organized, and how do atomic properties change across it?
The periodic table and periodic trends: describe the organization of the periodic table and the trends in atomic radius, ionization energy, electronegativity and reactivity across periods and down groups.
A focused Virginia SOL Chemistry answer on the periodic table under CH.2: how it is organized into groups, periods, metals, nonmetals and metalloids, and the trends in atomic radius, ionization energy, electronegativity and reactivity and why each runs the way it does.
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What this topic is asking
Standard CH.2 asks you to read the periodic table as an organizational tool and to describe the periodic trends in atomic radius, ionization energy, electronegativity and reactivity. The SOL tests both the direction of each trend and a short explanation in terms of atomic structure, and the periodic table provided on the test lets you locate any element to apply the trends.
How the periodic table is organized
Several families are named: the alkali metals (Group 1, very reactive metals), the alkaline earth metals (Group 2), the halogens (Group 17, very reactive nonmetals), and the noble gases (Group 18, full outer levels and almost unreactive). The repeating pattern of properties is "periodicity", which is why the table works as a predictive tool.
Atomic radius
The largest atoms are at the bottom left (such as francium and cesium) and the smallest are at the top right (excluding the noble gases). When an atom loses electrons to form a cation it gets smaller; when it gains electrons to form an anion it gets larger.
Ionization energy
Ionization energy increases across a period (a higher nuclear charge in a smaller atom holds electrons tightly) and decreases down a group (the outer electron is farther away and shielded, so it is easier to remove). This is why metals at the bottom left lose electrons readily while nonmetals at the top right resist losing them.
Electronegativity and reactivity
Electronegativity is the tendency of an atom to attract shared electrons in a bond. It increases across a period and decreases down a group, so fluorine (top right) is the most electronegative element. This trend feeds straight into bonding: a large electronegativity difference gives an ionic bond and a small one a covalent bond.
Reactivity follows the ease of gaining or losing electrons. Metals become more reactive down a group (lower ionization energy, lose electrons more easily), so cesium is more reactive than lithium. Nonmetals become more reactive up a group (higher electronegativity), so fluorine is the most reactive halogen.
Try this
Q1. Which element has the higher first ionization energy, sodium or chlorine? Explain briefly. [1 point]
- Cue. Chlorine; it is farther right with a greater nuclear charge and smaller radius, so it holds its electrons more tightly.
Q2. State the trend in metallic reactivity down Group 1 from lithium to cesium. [1 point]
- Cue. Reactivity increases down the group, because ionization energy falls and the outer electron is lost more easily.
Exam-style practice questions
Practice questions written in the style of VDOE exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
SOL (multiple choice)1 marksAs you move from left to right across Period 2 from lithium to fluorine, the atomic radius generally (A) increases (B) decreases (C) stays the same (D) increases then decreasesShow worked answer →
The answer is (B) decreases.
Across a period the number of protons (nuclear charge) increases while electrons are added to the same outermost energy level. The stronger nuclear pull on the same shell draws the electron cloud inward, so the atomic radius decreases from left to right. The periodic table groups elements so this trend is consistent across each period.
The trap is assuming more electrons makes a bigger atom; across a period the rising nuclear charge dominates and the atom shrinks.
SOL (tech-enhanced, fill in blank)3 marksConsider the Group 17 elements fluorine, chlorine and bromine. (a) State the trend in atomic radius down the group. (b) State the trend in electronegativity down the group. (c) Explain the electronegativity trend in terms of atomic structure.Show worked answer →
A 3-point group-trend item.
(a) Atomic radius (1 point): increases down the group (fluorine smallest, bromine largest).
(b) Electronegativity (1 point): decreases down the group.
(c) Explanation (1 point): going down the group, atoms gain energy levels, so the outer electrons are farther from the nucleus and more shielded by inner electrons; the nucleus attracts bonding electrons less strongly, lowering electronegativity.
Markers reward the correct directions of both trends and an explanation linking larger size and greater shielding to weaker attraction.
Related dot points
- Structure of the atom: describe protons, neutrons and electrons, atomic number and mass number, and the historical development of the atomic model from Dalton to the modern view.
A focused Virginia SOL Chemistry answer on standard CH.2: the subatomic particles, atomic number and mass number, how they define an element and its ions, and the development of the atomic model from Dalton, Thomson and Rutherford to Bohr and the modern model.
- Electron configuration and energy levels: describe how electrons occupy energy levels, write electron configurations, identify valence electrons, and relate ground and excited states to spectra.
A focused Virginia SOL Chemistry answer on electron arrangement under CH.2: energy levels and sublevels, writing electron configurations, counting valence electrons, and the difference between ground state and excited state and how it produces line spectra.
- Isotopes and average atomic mass: define isotopes, write nuclide notation, and calculate the weighted average atomic mass of an element from its isotopes.
A focused Virginia SOL Chemistry answer on isotopes under CH.2: what isotopes are, how to read nuclide notation, and how to calculate the weighted average atomic mass of an element from the masses and natural abundances of its isotopes.
- Types of chemical bonds: explain ionic, covalent and metallic bonding in terms of valence electrons and electronegativity, and predict bond type from the periodic table.
A focused Virginia SOL Chemistry answer on bonding under CH.3: why atoms bond to reach a stable octet, how ionic, covalent and metallic bonds form, and how to predict the bond type from electronegativity difference and position on the periodic table.
- Polarity and intermolecular forces: determine molecular polarity from shape and bond polarity, and compare dispersion, dipole-dipole and hydrogen-bonding forces and their effect on properties.
A focused Virginia SOL Chemistry answer on polarity under CH.3: how bond polarity and molecular shape combine to make a molecule polar or nonpolar, the three intermolecular forces (dispersion, dipole-dipole, hydrogen bonding), and how they set boiling and melting points and solubility.
Sources & how we know this
- 2018 Science Standards of Learning - Chemistry — Virginia Department of Education (2018)
- Chemistry Curriculum Framework — Virginia Department of Education (2018)