What are the charges and locations of the three subatomic particles?

The proton has a charge of $1+$ and sits in the nucleus; the neutron has no charge and also sits in the nucleus; the electron has a charge of $1-$ and occupies the space outside the nucleus. Protons and neutrons each have a relative mass of about $1$, while the electron's mass is negligible, so the nucleus holds nearly all the atom's mass. In a neutral atom the number of protons equals the number of electrons.

How do you calculate average atomic mass on the Regents exam?

Multiply each isotope's mass by its natural abundance written as a decimal, then add the products: average $= (m_1 \times f_1) + (m_2 \times f_2) + \dots$. For example, chlorine is about $75.8\%$ chlorine-35 and $24.2\%$ chlorine-37, giving roughly $35.45$. The answer always lies between the lightest and heaviest isotope masses and closer to the more abundant isotope, which is why the periodic-table value is a decimal.

How do you tell a ground state from an excited state in New York configuration?

First add the numbers in the configuration; the total must equal the atomic number, or it is a different element. Then check whether the lowest energy levels are filled first. The ground state fills the lowest levels first (the configuration shown on the Periodic Table). An excited state has the same number of electrons but with one promoted to a higher level, leaving a lower level not completely filled, for example neon as $2\text{-}7\text{-}1$ rather than $2\text{-}8$.

What are the main periodic trends and why do they happen?

Across a period, atomic radius decreases while ionization energy and electronegativity increase, because nuclear charge rises as electrons fill the same level. Down a group, atomic radius increases while ionization energy and electronegativity decrease, because each element adds an energy level so outer electrons are farther away and more shielded. Metallic character is the opposite of ionization energy: it increases down a group and decreases across a period. Table S lists the actual values.

Which Reference Tables matter for this topic?

The Periodic Table itself gives atomic number, atomic mass and the ground-state electron configuration of each element. Table S (Properties of Selected Elements) lists atomic radius, electronegativity and first ionization energy, which trend questions and bond-polarity questions draw on. Knowing where these are and reading them quickly turns several marks into simple look-ups.

New YorkChemistry

Regents Chemistry atomic structure and the periodic table: a complete skills guide to particles, isotopes, configuration and trends

A deep-dive Regents Chemistry guide to atomic structure and the periodic table: the subatomic particles, isotopes and weighted average atomic mass, New York electron configuration with ground and excited states, the organization of the periodic table, and the trends in radius, ionization energy and electronegativity, plus the Reference Tables and exam technique.

Generated by Claude Opus 4.816 min read3.1Updated 2026-06-13

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Jump to a section

Why this is the foundation of Regents Chemistry
The three particles and how to count them
Isotopes and weighted average mass
Electron configuration, ground and excited states
Organizing the periodic table
The four trends
Ions and nuclide notation
Check your knowledge

Why this is the foundation of Regents Chemistry

Atomic structure and the periodic table are the first unit of the Physical Setting/Chemistry Core Curriculum, and everything after them depends on getting them automatic. The particles in an atom set up bonding and reactions; isotopes explain the masses you use in stoichiometry; electron configuration explains valence electrons and the periodic trends; and the periodic table is the single most-used page in the Reference Tables. This guide ties together the matching dot-point pages, each with its own Regents-format practice: subatomic particles and atomic structure, isotopes and average atomic mass, electron configuration and energy levels, the periodic table and its organization, periodic trends, and ions and nuclide notation.

The three particles and how to count them

An atom has a dense nucleus of protons ( $1+$ ) and neutrons (no charge), surrounded by electrons ( $1-$ ). The atomic number ( $Z$ ) is the proton count and defines the element; the mass number ( $A$ ) is protons plus neutrons. So neutrons $= A - Z$ , and a neutral atom has electrons equal to $Z$ . The nucleus carries almost all the mass because the electron's mass is negligible, and the atom is mostly empty space, which is what Rutherford's gold-foil experiment showed.

Isotopes and weighted average mass

Isotopes are atoms of the same element with different numbers of neutrons (same $Z$ , different $A$ ). Because the periodic-table mass is a weighted average over an element's isotopes, it is usually a decimal. When a question gives isotope masses and abundances, compute:

\text{average atomic mass} = (m_1 \times f_1) + (m_2 \times f_2) + \dots

with abundances as decimals that sum to $1$ . The answer lands between the isotope masses, closer to the more common one. A plain average is wrong unless the isotopes are equally abundant.

Electron configuration, ground and excited states

New York writes configurations as electrons per energy level, lowest first: sodium is $2\text{-}8\text{-}1$ . The ground state fills the lowest levels first (the form shown on the Periodic Table). An excited state has the same electron count but one electron promoted, so a lower level is not full ( $2\text{-}7\text{-}1$ for neon). Electrons absorb a specific amount of energy to jump up and release that energy as a photon when they fall back, producing each element's unique bright-line spectrum. The outermost-level electrons are the valence electrons, the ones that do the bonding.

Organizing the periodic table

Elements are arranged by increasing atomic number into periods (rows, same number of occupied levels) and groups (columns, same number of valence electrons). The periodic law says properties repeat regularly with atomic number. Metals (left and center) conduct, are malleable and form cations; nonmetals (upper right) are poor conductors and form anions; metalloids lie along the staircase. The named families are the alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17) and noble gases (Group 18). Bromine and mercury are the only liquid elements at room temperature.

The four trends

Property	Across a period (left to right)	Down a group (top to bottom)
Atomic radius	decreases	increases
First ionization energy	increases	decreases
Electronegativity	increases	decreases
Metallic character	decreases	increases

The reason is structural. Across a period the nuclear charge rises while electrons fill the same level, pulling the atom in and holding electrons tighter. Down a group, added energy levels put outer electrons farther out and shield them, so they are easier to remove and attract bonding electrons less. Table S lists radius, electronegativity and ionization energy for named elements.

Ions and nuclide notation

An ion forms when an atom loses electrons (a cation, positive) or gains them (an anion, negative). Protons and neutrons never change, so the element stays the same. Nuclide notation $^{A}_{Z}\text{X}^{\,c}$ gives protons $= Z$ , neutrons $= A - Z$ , and electrons $= Z - c$ . Cations are smaller than their atoms; anions are larger.

Atoms have a nucleus of protons ( $1+$ ) and neutrons (no charge) with electrons ( $1-$ ) outside; atomic number $Z$ is the proton count and mass number $A$ is protons plus neutrons, so neutrons $= A - Z$ . Isotopes differ only in neutrons, and an element's atomic mass is the weighted average of its isotopes. New York configurations list electrons per level, lowest first; the ground state fills lowest levels first and an excited state promotes an electron while keeping the same total. The periodic table runs by atomic number into periods and groups, with metals on the left and nonmetals at the upper right. Across a period radius falls while ionization energy and electronegativity rise; down a group the reverse holds. Ions form by losing or gaining electrons (protons fixed), and nuclide notation $^{A}_{Z}\text{X}^{\,c}$ gives all three particle counts.

Check your knowledge

Attempt these under timed conditions, then check the solutions.

State the charge and location of each of the three subatomic particles. (3 marks)
An atom has $17$ protons and $18$ neutrons. State its mass number and identify the element. (2 marks)
An element has isotopes of mass $10.0$ ( $20\%$ ) and $11.0$ ( $80\%$ ). Show the setup and calculate the average atomic mass. (2 marks)
Decide whether $2\text{-}8\text{-}4$ and $2\text{-}7\text{-}5$ are ground or excited states, and name the element each represents. (2 marks)
State the trend in atomic radius across Period 2 and explain it. (2 marks)
State the number of electrons in an $\text{O}^{2-}$ ion (oxygen has $8$ protons). (1 mark)
From the nuclide $^{40}_{20}\text{Ca}^{2+}$ , state the protons, neutrons and electrons. (3 marks)

Solutions

Q1: Proton: $1+$ , in the nucleus. Neutron: no charge, in the nucleus. Electron: $1-$ , outside the nucleus.
Q2: Mass number $= 17 + 18 = 35$ ; atomic number $17$ is chlorine.
Q3: $(10.0)(0.20) + (11.0)(0.80) = 2.0 + 8.8 = 10.8$ .
Q4: $2\text{-}8\text{-}4$ totals $14$ (silicon) and fills lowest first, so it is the ground state. $2\text{-}7\text{-}5$ also totals $14$ (silicon) but the second level is not full while the third is occupied, so it is an excited state of silicon.
Q5: Atomic radius decreases across Period 2 because nuclear charge increases while electrons stay in the same energy level, pulling the cloud inward.
Q6: $8 + 2 = 10$ electrons (oxygen gains two electrons to form $\text{O}^{2-}$ ).
Q7: Protons $= 20$ ; neutrons $= 40 - 20 = 20$ ; electrons $= 20 - 2 = 18$ .

Sources & how we know this

Physical Setting/Chemistry Core Curriculum — New York State Education Department (2002)
Reference Tables for Physical Setting/Chemistry, 2011 Edition — New York State Education Department (2011)