atoms

Matter and Elements: Matter consists of elements in pure form and in compounds. An element cannot be broken down chemically, while a compound consists of two or more elements.
Essential Elements for Life: About 20-25% of natural elements are essential; 96% of living matter is made of C, H, O, N; remaining 4% includes Ca, P, K, S; trace elements are required in minute quantities.

Atom Definition: Smallest unit of matter retaining element properties. Composed of neutrons, protons, and electrons.
Atomic Number and Mass: Atomic number = protons; mass number = protons + neutrons. Atomic mass approximated by mass number. Neutrons = mass number - atomic number.
Isotopes: Atoms of the same element differing in neutron number. Radioactive isotopes decay and have applications in biological research (e.g., dating fossils, medical diagnosis).

Covalent Bonds: Form through sharing valence electrons; can be single (one pair) or double (two pairs). The atom's bonding capacity is called valence.
Polar and Nonpolar Bonds: Nonpolar bonds share electrons equally; polar bonds do not, leading to partial charges.
Ionic Bonds: Form when atoms transfer electrons, creating cations (positive) and anions (negative) that attract each other.
Weak Interactions: Include hydrogen bonds and Van der Waals interactions, crucial for biological functions.

Definition: Making/breaking of chemical bonds; reactants transform into products.
Photosynthesis Example: Conversion of carbon dioxide and water to glucose and oxygen, illustrating a key chemical reaction.
Reversibility and Equilibrium: Reactions are reversible; equilibrium is reached when reactant and product rates are equal, and concentrations remain constant.

Matter and Elements: Matter consists of elements in pure and compound forms. An element is a fundamental substance that cannot be broken down through chemical methods, while a compound is formed by the chemical combination of two or more elements. Each element is represented by a unique atomic number, which corresponds to the number of protons in its nucleus.
Essential Elements for Life: Approximately 20-25% of the natural elements are essential for life. The majority (96%) of living matter is composed of four primary elements: Carbon (C), Hydrogen (H), Oxygen (O), and Nitrogen (N). The remaining 4% includes elements such as Calcium (Ca), Phosphorus (P), Potassium (K), and Sulfur (S). There are also trace elements that are essential in minute amounts, playing critical roles in biological processes.

Atom Definition: An atom is the smallest unit of matter that retains all the properties of an element. It is made up of three types of subatomic particles: neutrons (neutral charge), protons (positive charge), and electrons (negative charge). The number of protons defines the element itself.
Atomic Number and Mass: The atomic number of an element is equal to the number of protons in its nucleus. The mass number is the sum of protons and neutrons within an atom. Hence, the atomic mass is often approximated by the mass number since it accounts for the majority of the atom's mass. The number of neutrons can be found using the formula: Neutrons = Mass Number - Atomic Number.
Isotopes: Isotopes are variants of the same element that differ in the number of neutrons. They can exhibit different physical properties and some isotopes are radioactive. Radioactive isotopes decay over time and are useful in various applications, including radiometric dating and medical diagnostics, such as imaging and cancer treatment.

Covalent Bonds: These bonds form when two atoms share valence electrons. Depending on the number of electron pairs shared, covalent bonds can be classified as single (one pair of electrons) or double (two pairs of electrons). The bonding capacity of an atom is determined by its valence, which indicates how many electrons are available for bonding.
Polar and Nonpolar Bonds: In nonpolar covalent bonds, electrons are shared equally between atoms, which leads to a balanced charge distribution. In polar covalent bonds, there is an unequally shared electron distribution, creating partial charges on the atoms. This property is crucial in influencing molecular interactions and solubility.
Ionic Bonds: These bonds result from the complete transfer of electrons between atoms, leading to the formation of charged ions: cations (positively charged) and anions (negatively charged). The electrostatic attraction between the oppositely charged ions holds them together.
Weak Interactions: Weak interactions, such as hydrogen bonds and Van der Waals forces, are less strong than covalent and ionic bonds but are essential for biological functions. Hydrogen bonds, for instance, play a vital role in the structure of water and the secondary structure of proteins.

Definition: A chemical reaction involves the making and breaking of chemical bonds, leading to the transformation of reactants into products. During this process, the atoms involved are rearranged, and new substances are formed.
Photosynthesis Example: A classic example of a chemical reaction is photosynthesis, where carbon dioxide ( ext{CO}2) and water ( ext{H}_2 ext{O}) are converted into glucose ( ext{C}_6 ext{H}{12} ext{O}6) and oxygen ( ext{O}_2) in the presence of sunlight. This reaction can be represented as: $ext{6CO}_2 + 6 ext{H}_2 ext{O} ightarrow ext{C}_6 ext{H}</em>{12} ext{O}_6 + 6 ext{O}_2$
Reversibility and Equilibrium: Many chemical reactions are reversible, meaning the products can be converted back into reactants. Equilibrium occurs when the forward and reverse reaction rates are equal, resulting in constant concentrations of reactants and products over time.