Unit 2

Upon completion of this section, you should be able to:

Matter is defined as anything that takes up space and has mass. It exists in several states: solid, liquid, gas, or plasma (a state of ionized gas).
Living organisms and nonliving things (like air) are composed of matter.

All matter comprises elements, which are substances that cannot be broken down by ordinary chemical means.
Elements serve as the building blocks of matter. Naturally occurring elements are vital for biological functions.
Four elements—carbon, hydrogen, nitrogen, and oxygen—comprise approximately 96% of body weight in most organisms. Other essential elements include phosphorus, calcium, and sulfur, while trace elements include minerals like zinc and chromium.

Atomic Theory: Elements consist of tiny particles called atoms, each represented by an atomic symbol (e.g., H for hydrogen, Na for sodium).
Atoms are primarily made of three types of subatomic particles:
- Neutrons: No electrical charge.
- Protons: Positive charge.
- Electrons: Negative charge.
Protons and neutrons reside in the nucleus (center of the atom), while electrons orbit the nucleus. Most of the atom is empty space.
Analogy: If an atom were the size of a baseball stadium, the nucleus would be like a gumball, and electrons would be tiny specks moving around in the upper stands.

The mass number of an atom equals the sum of its protons and neutrons. Protons and neutrons each have a mass of one atomic mass unit (amu). Electrons are negligible in mass for most calculations.
Atomic number: Represents the number of protons in an atom, determining an element's identity. For neutral atoms, the number of electrons equals the number of protons.
To find the number of neutrons, subtract the number of protons from the atomic mass, rounded to the nearest whole number.

The Periodic Table displays elements based on their recurring chemical and physical properties.
Structure: Atomic numbers are written above the symbols, while atomic masses are provided below.
Example: Carbon (C) has:
- Atomic Number: 6
- Atomic Mass: 12.01
Elements in group 8 are called noble gases, which rarely react with other atoms due to complete outer shells.

Isotopes: Atoms of the same element with varying numbers of neutrons, hence different mass numbers.
Some isotopes are unstable and radioactive, meaning they can decay and emit radiation, detectable by devices like a Geiger counter. Not all isotopes are radioactive.

Radioactivity is utilized for electricity, food safety, and medical advancements. It's essential for understanding biological processes at the cellular level.

Radioactive isotopes can act as tracers to detect changes at the molecular level or to eliminate harmful cells. They behave similarly to stable isotopes during chemical reactions.
Example: Radioactive iodine (iodine-131) studies thyroid function. Positron-emission tomography (PET) scans utilize radiolabeled glucose to assess metabolic activity in tissues.

The model of electron arrangement likens electrons to a ceiling fan; when in motion, they blend into a blur. Their precise location is indeterminate but governed by energy levels called electron shells. More energy is required to orbit further from the nucleus.
Electron Shells: Electrons fill shells from the inside out—first shell (2 electrons), second shell (8 electrons), third shell and beyond (up to 18 electrons with increasing distance).
Octet Rule: Atoms aim to have eight electrons in their outer shell for stability, except those with fewer than two shells (which complete with two electrons).

Atoms can bond to form molecules (groups of atoms bonded together) or compounds (molecules with multiple element types).

Ionic bonds form from the attraction between oppositely charged ions resulting from electron transfer.
Example: In sodium chloride (NaCl): Sodium (Na) donates an electron, becoming Na+; Chlorine (Cl) accepts the electron, becoming Cl−. Together they form NaCl.

Covalent bonds occur when atoms share electrons to achieve filled outer shells.
Example: Hydrogen gas (H₂) involves shared electrons, illustrated by the structural formula H-H.
Atoms can share multiple pairs of electrons, forming double (like O₂) or triple bonds (like N₂).

Differences in ionic vs covalent bonding include the method of electron arrangement: ionic involves transfer and resulting charge, while covalent involves sharing.

Methane (CH₄): Carbon forms tetravalent bonds with four hydrogen atoms, with bonds represented in various models (ball-and-stick, space-filling).