Introduction to the Atom: Protons, Neutrons, Electrons, and Ions

Introduction to the Atom

Lecture Topics:
1. Protons
2. Neutrons
3. Electrons
4. Intro to Ions
Reminders:
- Discussion on Friday.
- GSS (Group Study Session) starts next week; announcements will be posted.
- HW #1 opens September 4th on MC (due September 18th by 11:59 pm).
Humorous Aside: "DO NOT TRUST ATOMS, THEY MAKE UP EVERYTHING."

Accompanying Readings and Practice Problems

Accompanying Readings: Section 2.6
Recommended Practice Problems: 2.41, 2.49, 2.61, 2.63, 2.65, 2.67

What is Chemistry and Matter?

What is Chemistry? The study of matter.
What is Matter? Anything that has mass and takes up space.
What does it mean when we say matter is particulate? Matter is made up of "stuff."
What is the "stuff"? Atoms.

Dalton's Theory of the Atom (1803)

Dalton's theory, proposed in 1803, laid foundational concepts for the atom, although three out of its four postulates have since been disproven.

Elements are made up of tiny, indestructible particles called atoms.
- False: Atoms are not indestructible; they are composed of subatomic particles (protons, neutrons, electrons).
Atoms of one element cannot change into other atoms of another element – they can only change the way they are bound with other atoms.
- False: Atoms can change into other elements through processes like radioactive decay.
All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.
- False: Isotopes exist, which are atoms of the same element with different masses due to varying numbers of neutrons.
Atoms combine in whole number ratios to form compounds.
- True: This postulate remains a fundamental principle of chemistry, describing how elements combine to form molecules (e.g., $ext{1 atom C}$ with $ext{1 atom O}$ for carbon monoxide, $ext{1 atom C}$ with $ext{2 atoms O}$ for carbon dioxide).

Subatomic Particles

Atoms are composed of three primary subatomic particles: protons, neutrons, and electrons.

Particle	Mass (kg)	Mass (amu)	Charge (relative)	Charge (C)
Proton	$1.67262 imes 10^{-27}$	$1.00727$	$1+$	$+1.60218 imes 10^{-19}$
Neutron	$1.67493 imes 10^{-27}$	$1.00866$	$0$	$0$
Electron	$0.00091 imes 10^{-27}$	$0.00055$	$1-$	$-1.60218 imes 10^{-19}$

Relative Masses: Protons and neutrons have approximately equal masses, with neutrons being slightly heavier. Electrons have a significantly smaller mass (approximately $rac{1}{1836}$ that of a proton). Therefore, the mass of an atom is mainly determined by its protons and neutrons.
Roles of Subatomic Particles:
- Proton (p^+$): Identifies an atom. The number of protons determines the element (e.g., an atom with 6 $protons is Carbon). It carries a$ 1+ $relative charge.</li><li>Electron ($ e^-$): Determines the atom's charge and reactivity. Atoms can gain or lose electrons. It carries a $1-$ relative charge. The mass change due to gaining or losing electrons is negligible.
- Neutron (n^ ext{o}$): Primarily contributes to the stability of the atom's nucleus. It carries no charge.

The Periodic Table and Atomic Information

The Periodic Table organizes elements based on their atomic number and electron configurations, revealing periodic trends in properties.

Groups (Columns): Elements in the same group (vertical column) tend to have similar chemical reactivity.
Periods (Rows): Elements in the same period (horizontal row) show trends in properties across the row.

Each element entry on the periodic table typically displays:

Atomic Number (Z $): Located at the top of an element's symbol (e.g.,$ 6 $for Carbon). This number represents the number of protons in the nucleus of an atom and uniquely identifies the element.</li><li>Chemical Symbol: A one or two-letter abbreviation for the element (e.g.,$ ext{C} $for Carbon).</li><li>Average Atomic Mass: Found below the chemical symbol (e.g.,$ 12.01 $for Carbon). This value is the weighted average of the masses of all naturally occurring isotopes of that element, expressed in atomic mass units (amu). This mass also represents the molar mass when expressed in grams per mole ($ ext{g/mol} $), meaning$ 1 $mole of atoms of that element weighs this amount in grams.</li></ul><h4 id="af3ca1d6-550f-4e17-9a8c-91898e87f9de" data-toc-id="af3ca1d6-550f-4e17-9a8c-91898e87f9de" collapsed="false" seolevelmigrated="true">Neutrons ($ n^ ext{o} $) and Isotopes</h4><ul><li>Role in Stability: Neutrons play a crucial role in the stability of the atomic nucleus. For lighter elements, a stable atom often has a neutron-to-proton ratio ($ rac{N}{Z} $) close to$ 1 $. As the atomic number ($ Z $) increases, the number of neutrons ($ N $) generally needs to increase faster than the number of protons to maintain stability (the N/Z curve typically goes above$ 1 $for heavier stable nuclei).</li><li>Isotopes Defined: Isotopes are two atoms of the same element (meaning they have the same number of protons) but have different numbers of neutrons. This difference in neutron count leads to different atomic masses for isotopes of the same element.</li><li>Examples of Carbon Isotopes:<ul><li>$ ext{C-12} $: The most abundant isotope of carbon.</li><li>$ ext{C-13} $: Has one more neutron than$ ext{C-12} $, with an approximate natural abundance of$ 1.1 ext{%} $.</li><li>$ ext{C-14} $: Has two more neutrons than$ ext{C-12} $and is radioactive, with negligible natural abundance in stable forms (used in carbon dating).</li></ul></li></ul><h4 id="f99c0559-8403-453c-8ddf-ba44dee6fa00" data-toc-id="f99c0559-8403-453c-8ddf-ba44dee6fa00" collapsed="false" seolevelmigrated="true">Electrons and Charged Atoms (Ions)</h4><ul><li>Neutral Atoms: In a neutral atom, the number of electrons is equal to the number of protons ($ Z $). This balances the positive charge of the protons with the negative charge of the electrons, resulting in an overall net charge of zero.</li><li>Ions: During chemical changes, atoms can either lose or gain electrons to achieve a more stable electron configuration (often mimicking noble gases). When an atom gains or loses electrons, it becomes a charged particle called an ion.</li><li>Cations:<ul><li>Definition: Positively charged ions.</li><li>Formation: Formed when an atom loses electrons.</li><li>Example: Metal elements, such as Sodium ($ ext{Na} $), tend to form cations. The reaction for forming a sodium cation is$ ext{Na}
ightarrow ext{Na}^+ + ext{e}^- $. (Mnemonic: "Cats have paws," referring to the positive charge of cations).</li></ul></li><li>Anions:<ul><li>Definition: Negatively charged ions.</li><li>Formation: Formed when an atom gains electrons.</li><li>Example: Nonmetal elements, such as Fluorine ($ ext{F} $), tend to form anions. The reaction for forming a fluoride anion is$ ext{F} + ext{e}^-
ightarrow ext{F}^- $.</li></ul></li></ul><h4 id="556b9b42-ab58-4d8d-b9f3-0dce74697440" data-toc-id="556b9b42-ab58-4d8d-b9f3-0dce74697440" collapsed="false" seolevelmigrated="true">Examples of Ion Formation</h4>Let's consider how common elements form ions:<ol><li>Neon ($ ext{Ne} $), Atomic Number ($ Z = 10 $):<ul><li>In a neutral$ ext{Ne} $atom:$ 10 $protons ($ p^+$), $10$ electrons (e^-$).
Neon is a noble gas, which means it already has a stable electron configuration and does not readily gain or lose electrons.

Fluorine ( ext{F} $), Atomic Number ($ Z = 9 $):<ul><li>In a neutral$ ext{F} $atom:$ 9 $protons ($ p^+$), $9$ electrons (e^-$).

To achieve a stable configuration like Neon (its nearest noble gas), Fluorine will gain one electron.

As an ion ( ext{F}^- $):$ 9 $protons ($ p^+$), $10$ electrons (e^-$). This is an anion.

Magnesium ( ext{Mg} $), Atomic Number ($ Z = 12 $):<ul><li>In a neutral$ ext{Mg} $atom:$ 12 $protons ($ p^+$), $12$ electrons (e^-$).

To achieve a stable configuration like Neon (its nearest noble gas), Magnesium will lose two electrons.

As an ion ( ext{Mg}^{2+} $):$ 12 $protons ($ p^+$), $10$ electrons ($$e^-$). This is a cation.