Matter, Elements, and Compounds

Matter, Elements, and Compounds

• Matter is anything that occupies space and has mass and can be found in solid, liquid, or gas form
• Elements are substances that cannot be broken down further by chemical means
• There are 92 naturally occurring elements like gold, copper, carbon and oxygen
• Compounds contain two or more elements in a fixed ratio, such as table salt (NaCl), made of sodium (Na) and chlorine (Cl)

Essential Elements for Life

• Six elements make up about 99% of the human body: oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus
• The first four are key components of proteins, carbohydrates, and lipids
• Calcium and phosphorus are major components of bones and teeth

Trace Elements

• Trace elements make up less than 0.01% of human body weight
• Trace elements include boron, chromium, cobalt, copper, fluorine, iodine, iron, manganese, molybdenum, selenium, silicon, tin, vanadium, and zinc
• Iron (Fe) is essential for oxygen transport in blood as part of hemoglobin
• Zinc (Zn) is important for enzyme function, immune system health, and DNA synthesis
• Iodine (I) is required to produce thyroid hormones that regulate metabolism
• Copper (Cu) plays a role in energy production and the formation of connective tissues
• Selenium (Se) acts as an antioxidant to protect cells from damage
• Manganese (Mn) is involved in bone formation and enzyme activation

Atomic Structure

• An atom, derived from the Greek word for "indivisible," is the smallest unit retaining element properties
• The nucleus contains protons (positive charge) and neutrons (neutral charge)
• Electrons are negatively charged subatomic particles
• Neutrons are electrically neutral

Atomic Number vs. Atomic Mass

• The atomic number defines an element by its unique number of protons; helium has an atomic number of 2 because it has 2 protons
• Atoms typically have an equal number of protons and electrons, resulting in a neutral charge
• The mass number is the sum of protons and neutrons; helium's mass number is 4
• Atomic mass is approximately equal to the mass number, measured in daltons

Stable vs. Radioactive Isotopes

• Isotopes are variations of an element with the same atomic number but different neutron numbers
• Carbon has three isotopes: carbon-12 (6 neutrons), carbon-13 (7 neutrons), and carbon-14 (8 neutrons)
• Carbon-12 is the most common isotope, making up about 99% of naturally occurring carbon
• Stable isotopes, like carbon-12 and carbon-13, have nuclei that remain intact
• Radioactive isotopes, like carbon-14, decay spontaneously, emitting particles and energy -Radiation from radioactive isotopes can damage cellular molecules and pose risks to living organisms
• Carbon-14 is used in dating fossils and is employed in biological research and medicine
• Stable isotopes have a balanced ratio of protons to neutrons allowing the strong nuclear force to counteract forces
• Radioactive isotopes have an imbalance, making the nucleus unstable

Electron Shells and Chemical Bonds

• Electron shells are regions around the nucleus where electrons are likely to be found, organized into energy levels
• The innermost shell has the lowest energy, and the outermost (valence) shell has the highest energy
• Each shell holds a maximum number of electrons: 2 in the first shell, 8 in the second shell, and so on
• Atoms seek stability by filling their valence shell, leading them to gain, lose, or share electrons through chemical bonds

Ionic vs. Covalent Bonds

Ionic Bonds

• Electrons are transferred, creating charged ions (positive and negative)
• These bonds are strong due to electrostatic attraction but can weaken in water
• An example is sodium chloride (NaCl), where sodium gives up an electron to chlorine
• Ionic bonds typically form crystalline solids that dissolve in water and conduct electricity

Covalent Bonds

• Atoms share pairs of electrons to fill their valence shells
• These bonds are generally stronger than ionic bonds
• Water (H₂O) is an example, where oxygen shares electrons with two hydrogens
• Covalent bonds usually form molecules with variable solubility that often do not conduct electricity

Key Differences

• Ionic bonds transfer electrons and form lattice structures
• Covalent bonds share electrons and form discrete molecules
• Ionic compounds are hard and brittle, while covalent compounds vary in texture and state

Key Similarities

• Both ionic and covalent bonds aim to achieve stable electron configurations
• Both play a key role in matter's structure and properties

Electronegativity

• A molecule consists of two or more atoms held together by covalent bonds
• Electronegativity measures an atom's attraction for shared electrons
• In a bond between identical atoms, electrons are shared equally, forming nonpolar covalent bonds
• Unequal sharing happens when electronegativity differs and forms polar covalent bonds
• Oxygen, fluorine, and nitrogen are highly electronegative

Polar vs. Non-Polar Covalent Bonds

Polar Covalent Bonds

• Electrons are shared unequally due to differences in electronegativity
• This results in uneven electron distribution, creating partial charges
• Water (H₂O) is an example, with oxygen being more electronegative than hydrogen
• Polar molecules are usually soluble in water and have higher boiling/melting points

Non-Polar Covalent Bonds

• Electrons are shared equally between atoms with similar electronegativity
• There are no partial charges due to even electron distribution
• Methane (CH₄) is an example, where carbon and hydrogen share electrons equally
• Non-polar molecules are typically insoluble in water but soluble in non-polar solvents and have lower boiling/melting points

Key Differences

• Polar bonds involve unequal sharing, while non-polar bonds involve equal sharing
• Polar bonds create partial charges, whereas non-polar bonds do not
• Polar molecules interact with each other, while non-polar molecules interact more readily with non-polar substances

Ions

• An ion is an atom or molecule with a net electrical charge, resulting from gaining or losing electrons

Electron Gain

• When an atom or molecule gains electrons, it becomes negatively charged and is called an anion; chloride (Cl⁻) is an anion

Electron Loss

• When an atom or molecule loses electrons, it becomes positively charged and is called a cation; sodium (Na⁺) is a cation
• Neutral atoms have an equal number of protons and electrons
• Changes in this balance create an ion with either a positive or negative charge

Salt

• Sodium chloride, i.e. salt, is an ionic compound that forms crystals
• Sodium (Na⁺) and chloride (Cl⁻) ions are always present in a 1:1 ratio in a sodium chloride crystal

Hydrogen Bonds in Water

• Water molecules have hydrogen atoms attached to oxygen atoms by polar covalent bonds
• This makes water a polar molecule with unequal charge distribution
• The oxygen end is slightly negative, and the hydrogen ends are slightly positive
• The partial positive charge allows hydrogen to be attracted to nearby partially negative atoms
• Hydrogen bonds occur when one atom in the attraction is always a hydrogen atom
• Each hydrogen atom in water can form a hydrogen bond with a nearby oxygen atom
• The negative oxygen pole can form bonds with two hydrogen atoms, allowing one water molecule to bond with up to four partners

Reactants and Products

• Reactants are the starting materials and are found on the left side of the equation
• Products are the materials resulting from the chemical reaction and are found on the right side of the equation

Balanced Chemical Equations

• Chemical reactions only rearrange matter rather than creating or destroying it
• Covalent bonds holding hydrogen atoms together in H₂ and oxygen atoms together in O₂ are broken, and new bonds are formed to yield the H₂O product molecules

Hydrogen Bonds, Adhesion, Cohesion, and Surface Tension of Water

Cohesion

• Refers to the attraction between water molecules
• Hydrogen bonds cause water molecules to stick together
• Cohesion explains water droplet formation and water's high boiling point

Adhesion

• Refers to water's ability to stick to other substances
• Hydrogen bonds form between water molecules and molecules on other surfaces
• As an example, water adheres to plant vessel walls allowing upward mobility during transpiration
• Adhesion is crucial for capillary action, allowing water to climb narrow tubes

Surface Tension

• Results from cohesion acting at the surface of water, which creates a film like layer that resists force
• Water striders are able to walk on water because of the water tension

Water as an Effective Heat Sink

• Water has a strong resistance to changes in temperature because of hydrogen bonding

Thermal Energy and Heat

• Thermal energy is the random movement of atoms and molecules
• Heat transfers thermal energy from a warmer body to a cooler one
• Temperature measures the intensity of heat

Slow Heating of Water

• Water heats up more slowly than metal because energy is used to break the hydrogen bonds, preventing the molecules from moving faster
• Conversely, cooling water molecules releases heat as hydrogen bonds form, allowing water to absorb and store heat effectively

Moderating Earth’s Temperatures

• Large bodies of water serve as a thermal buffer and stabilize air temperatures
• The Ocean's resistance to temperature change stabilizes the environment, creating conditions for marine life

Regulating Body Temperature

• Water makes up about 66% of the body
• Water's thermal properties help maintain stable internal conditions

Evaporative Cooling

• In evaporative cooling, high-energy molecules escape during evaporation, lowering energy and then cooling the surface
• This is seen in sweat and also cools the environment on a larger scale

Properties of Water

• Water is found as gas, liquid, and solid
• Water is less dense as a solid because hydrogen bonds form creating a three dimensional crystal with fewer molecules
• Floating ice insulates the water below, allowing aquatic life to survive

Solutions

Solute

• A substance being dissolved

Solvent

• A dissolving agent

Solution

• A solution is a liquid made of a uniform mixture

Solute vs. Solvent

• The solute is smaller and broken down
• The solvent is large with the capability to break down
• Sugar dissolving in tea is an example of solute and solvent

Water as a Versatile Solvent

• Water has positively charged hydrogen that attracts negative chloride ions
• Water's partial charge clings to positive sodium ions
• The water separates all the ions in a salt crystal

Acids vs. Bases

• An acid donates hydrogen ions , such as hydrochloric acid
• Acidic solutions have a higher concentration of H+ than OH-
• A base reduces hydrogen, such as sodium hydroxide
• Bases increase the OH- concentration

pH Scale

• pH measures how acidic or basic a solution is, 7 is neutral
• Values below 7 indicate acidity, while values above 7 indicate basicity
• There is a tenfold difference in the concentration

Buffers

• Biological fluids use buffers to maintain a constant pH by accepting/donating H+ ions