Every atom consists of a dense nucleus (protons + neutrons) surrounded by electrons in energy levels called shells or orbitals. The number of protons - the atomic number - uniquely identifies the element and never changes for a given element. The number of neutrons can vary, giving rise to isotopes.
The number of electrons in a neutral atom equals the number of protons, but electrons can be gained or lost to form ions.
Electrons do not orbit the nucleus like planets. They occupy regions of probability called orbitals, organized into shells (n = 1, 2, 3...) and subshells (s, p, d, f). The first shell holds 2 electrons, the second holds 8, and the pattern continues.
The arrangement of electrons in the outermost shell - the valence shell - determines how an element bonds and reacts.
Understanding Atomic Structure
Most of an atom's mass is packed into an extremely small central region. This explains why matter can feel solid even though atoms contain a great deal of empty space. The positive charge in the center attracts negative electrons.
That electrical attraction keeps electrons bound to the atom. Electrons closer to the center have lower energy and are held more tightly. Moving an electron farther out requires an input of energy.
When an excited electron returns to a lower energy state, it releases energy as light. Each element produces a characteristic set of light colours because its allowed electron energies are different. This is the basis of flame tests and the coloured lines seen in spectroscopy.
Neutrons have no electric charge, yet they have an important job in the nucleus. Protons repel one another because they are all positive. A stable nucleus needs the strong nuclear force to hold its particles together.
Neutrons add strong-force attraction without adding more electrical repulsion. For small atoms, stable forms often have similar numbers of protons and neutrons. Larger nuclei usually need extra neutrons.
Some neutron and proton combinations are unstable. These atoms change through radioactive decay and release particles or energy. Carbon fourteen is one example used to date once-living materials.
Students should separate the ideas of chemical change and nuclear change. Chemical reactions rearrange electrons. Nuclear changes alter the nucleus and can turn one element into another.
The mass shown for an element on the periodic table is usually not a whole number. It is an average based on the natural mixture of its isotopes. A sample with more of a heavier isotope has a slightly larger average atomic mass.
This differs from the mass number of one particular atom, which is a whole-number count of particles in its nucleus. Isotopes behave almost identically in ordinary chemical reactions because they have the same electron arrangement when neutral.
Their masses can still matter. Lighter and heavier isotopes move at slightly different rates, and scientists use this difference in medical scans, environmental studies, and tracing chemical pathways.
Electron transfer helps explain many familiar substances. When a metal atom loses an outer electron, its remaining positive charge attracts a negatively charged atom or group of atoms. This attraction forms ionic compounds such as sodium chloride.
Atoms can instead share outer electrons, forming covalent bonds in substances such as water and oxygen gas. A useful learning habit is to draw a simple particle count before deciding an atom's charge or mass. Count protons to identify the element.
Combine protons with neutrons for the mass number. Compare electrons with protons to find the overall charge. Keep these three counts separate, since mixing them up causes many early mistakes in atomic structure.
Key Facts
- Atomic number
- Mass number
- Isotopes: same number of protons, different number of neutrons - same element, different mass.
- Ion: atom that has gained or lost electrons. Positive ion (cation): lost electrons. Negative ion (anion): gained electrons.
- First shell: max 2 electrons. Second shell: max 8 electrons. Third shell: max 18 electrons.
- Valence electrons: electrons in the outermost occupied shell; determine bonding behavior.
Vocabulary
- Proton
- Positively charged particle in the nucleus; defines the element's atomic number.
- Neutron
- Neutral particle in the nucleus; contributes to atomic mass and distinguishes isotopes.
- Electron
- Negatively charged particle outside the nucleus; determines chemical behavior.
- Isotope
- Atoms of the same element with different numbers of neutrons (and therefore different mass numbers).
- Electron shell
- A region around the nucleus where electrons of a given energy level are found.
Common Mistakes to Avoid
- Confusing atomic number with atomic mass. Atomic number (whole number) counts protons; atomic mass (often a decimal) is the weighted average across isotopes.
- Thinking the mass number is always a whole number for a specific element. Individual atoms have integer mass numbers, but the periodic table shows average atomic mass across all natural isotopes.
- Forgetting that in an ion, the number of electrons changes but protons don't. Fe³⁺ still has 26 protons but only 23 electrons.
- Placing more than 8 electrons in the second shell in simple Bohr models. The Bohr model caps shells at 2, 8, 8, 18 - but the third shell holds more than 8 only after the 4s fills.
Practice Questions
- 1 An atom has 17 protons, 18 neutrons, and 17 electrons. Identify the element, write its atomic symbol, and state its mass number.
- 2 Carbon-12 and Carbon-14 are isotopes. How do they differ in atomic structure?
- 3 An ion has 11 protons and 10 electrons. What element is it, and what is the charge on the ion?