Chemistry: Matter and Elements Overview

Matter and Elements

• Matter occupies space and has mass.
• Matter exists as solids, liquids, and gases.
• All matter is composed of chemical elements.
• Elements are substances not broken down by ordinary chemical means.
• 92 naturally occurring elements exist.
• Chemists have created synthetic elements.
• Each element has a unique symbol (e.g., O for oxygen, Na for sodium).
• Each element has its own unique type of atom, distinct from the atoms of other elements.
• An atom is the smallest unit of matter that retains the properties of an element.
• Atoms are incredibly small, with about a million fitting across the period at the end of a sentence.

Atoms

• Physicists have identified over a hundred types of subatomic particles, but only three (protons, electrons, and neutrons) are relevant for atoms.
• Protons have a positive charge.
• Electrons have a negative charge.
• Neutrons have no charge.
• An atom's nucleus contains protons and neutrons tightly packed together.
• Electrons exist in a cloud around the nucleus; simpler models show electrons orbiting the nucleus.
• Electrons occupy electron shells at varying distances from the nucleus.

Atomic Structure and Properties

• Atoms of different elements are distinguished by their unique number of protons (atomic number).
• A helium atom with two protons has an atomic number of 2.
• Atoms typically have equal numbers of protons and electrons, resulting in a net zero charge.
• An atom's mass number is the sum of its protons and neutrons.
• For helium, the mass number is 4.
• Protons and neutrons have nearly identical mass (1 dalton).
• Electrons are much lighter and contribute negligibly to an atom's mass.
• An atom's atomic mass is roughly equal to its mass number in daltons.
• Atoms of the same element can have different mass numbers due to varying numbers of neutrons (isotopes).
• For example, carbon has three isotopes (carbon-12, carbon-13, and carbon-14), each with a different number of neutrons.
• Only electrons are directly involved in chemical activity.
• Helium's 2 electrons are in an inner circle around the nucleus. Atoms with more than 2 electrons have inner and outer electron shells.
• Elements are arranged in rows based on the number of electron shells.
• Electrons occupy orbitals within each shell, with up to 2 electrons per orbital.
• The valence shell is the outermost electron shell.
• An atom's chemical properties are primarily determined by the number of electrons in the valence shell.
• Atoms with incomplete outer shells tend to interact with other atoms to fill their valence shells.
• Reactive atoms, like hydrogen, carbon, nitrogen, and oxygen, have incomplete outer shells and readily interact with other atoms, forming bonds.
• Inert atoms, like helium, neon, argon, have complete outer shells, making them stable and less reactive.

Compounds

• Compounds are substances formed by two or more elements in a fixed ratio.
• Compounds have different properties than their constituent elements (e.g., NaCl, H₂O).
• Compounds are crucial building blocks of living organisms.

Elements in Living Organisms

• Humans need 25 essential elements for survival; plants need 17.
• Key elements (oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus) make up roughly 99% of the human body.
• These elements form proteins, carbohydrates, lipids, bones, and teeth.
• Other essential elements (potassium, sulfur, sodium, chlorine, magnesium) are needed for various biological functions like nerve signaling.

Trace Elements

• Trace elements are present in small amounts (less than 0.01% of body weight).
• Examples include iron, iodine, zinc, and others.
• Trace elements play vital roles in bodily functions, like energy processing and oxygen transport.
• Some trace elements are essential only for specific species.

Chemical Bonds

• Atoms with incomplete outer shells interact to fill their valence shells.

• Chemical bonds hold atoms close together.

• Ionic bonds form when one atom transfers an electron to another, creating an attraction.

• Covalent bonds form when atoms share electrons.

• The type of bond (and its strength) affects the properties of the resulting compound.

• A molecule consists of two or more atoms held together by covalent bonds.

• Electronegativity measures an atom's attraction for the shared electrons in a covalent bond.

• Nonpolar covalent bonds occur when atoms of the same element share electrons equally.

• Polar covalent bonds result when atoms with differing electronegativity share electrons unequally.

• The more electronegative atom has a partial negative charge.

• Atoms gain stability by completing their outer electron shells.

• Ionic bonds form when atoms with significantly different electronegativity transfer electrons, creating oppositely charged ions that attract each other.

• The transfer of electrons from sodium (Na) to chlorine (Cl) creates the ionic compound sodium chloride (NaCl), where sodium becomes positively charged (Na⁺) and chlorine becomes negatively charged (Cl⁻).

• Sodium chloride (NaCl), or table salt, is an ionic compound that forms crystals.

• The ratio of sodium (Na⁺) and chloride (Cl⁻) ions in a salt crystal is always 1:1.

• Ionic bonds are strong in dry salt crystals.

• When placed in water, the ionic bonds break, and the salt dissolves due to the attraction between the ions and water molecules.

• Many drugs are manufactured as salts to be stable when dry but dissolve easily in water.

Chemical Reactions

• Cells continuously rearrange molecules by breaking and forming chemical bonds.
• A chemical reaction changes reactants into products.
• The chemical equation 2 H₂ + O₂ → 2 H₂O shows hydrogen and oxygen combining to form water.
• Chemical reactions do not create or destroy matter; they rearrange it.
• Photosynthesis (6 CO₂ + 6 H₂O → C₆H₁₂O₆ + 6 O₂) is a crucial reaction in which plants use carbon dioxide and water with sunlight's energy to produce glucose and oxygen.
• Thousands of chemical reactions occur in the cells in the watery environment.

Water's Life-Supporting Properties

• Water's properties stem from its molecular structure and hydrogen bonding.
• Hydrogen bonds between water molecules are weak but numerous, creating cohesion, strong for water.
• Cohesion helps transport water and nutrients in trees (from roots to leaves).
• Evaporation at leaves creates a pulling force through cohesion.
• Adhesion (clinging of one substance to another) also aids water transport.
• Plant veins enhance water adhesion.
• Surface tension, related to cohesion, is relatively high for water.
• Hydrogen bonds create high surface tension, making water behave like a coated film.
• This surface tension supports some insects (e.g., water striders).
• Cohesion, adhesion, and water's high surface tension relate to structure and function in living organisms.
• This supports the concept of structure and function, where the plant vein structure supports water transport.
• Thermal energy is the energy associated with the random movement of atoms and molecules.
• When thermal energy transfers from a warmer to a cooler body, it is defined as heat.
• Temperature measures the intensity of heat, which is the average speed of molecules in a body of matter.
• Water heats up more slowly than metal due to hydrogen bonding, giving it a stronger resistance to temperature change compared to most other substances.
• Heat must be absorbed to break hydrogen bonds, and heat is released when hydrogen bonds form.
• To raise the temperature of water, hydrogen bonds between water molecules must be broken before the molecules can move faster.
• As a result, water absorbs a large amount of heat while warming up only slightly.
• Conversely, when water cools, water molecules slow down and more hydrogen bonds form, releasing a significant amount of heat.
• Earth's vast water supply helps moderate temperatures, keeping them within life-permitting limits.
• Oceans, lakes, and rivers absorb a large amount of heat from the sun during warm periods and release it gradually, warming the air.
• This is why coastal areas generally have milder climates than inland regions.
• Water's resistance to temperature change also stabilizes ocean temperatures, creating a favorable environment for marine life.
• Additionally, since water makes up about 66% of your body weight, it helps regulate your body temperature.
• When a substance evaporates (changes from a liquid to a gas), the surface of the remaining liquid cools down.
• This evaporative cooling happens because the molecules with the highest energy (the "hottest" ones) leave.
• Evaporative cooling helps prevent land-dwelling organisms from overheating.
• Evaporation from a plant's leaves keeps them cool in the sun, and sweating helps dissipate excess body heat in humans.
• The evaporation of surface waters cools tropical seas.
• Water is less dense as a solid (ice) than as a liquid due to hydrogen bonding. The ice crystal has fewer molecules than an equal volume of liquid water, making ice less dense, and allowing it to float.
• Floating ice insulates the water below, preventing it from freezing completely, allowing aquatic life to survive.
• In the Arctic, the ice serves as a hunting ground for polar bears.

Additional Information

• In living organisms, most strong chemical bonds are covalent.
• Weaker bonds, like ionic bonds, are crucial for cell function.
• Hydrogen bonds are a vital type of weak bond, especially important in water molecules.
• Water molecules have polar covalent bonds that arrange the positive and negative charges unevenly, which is crucial for hydrogen bonding.
• Hydrogen bonds occur when a hydrogen atom (covalently bonded to a highly electronegative atom like oxygen or nitrogen) attracts another electronegative atom.
• In water, each hydrogen can form a bond with a neighboring oxygen.
• A water molecule can form hydrogen bonds with up to four neighboring water molecules.
• Hydrogen bonds are crucial for protein structure, DNA structure, and the properties of water.

Terms to Know(added)

• Cohesion: Tendency of like molecules to stick together, especially high for water.
• Adhesion: Clinging of one substance to another; in organisms, structures enhance water adhesion.
• Solute: Substance dissolved.
• Solvent: Dissolving agent.
• Solution: Uniform mixture of two or more substances in a liquid.

Solutions and Solvents (Updated)

• When a solute (e.g., salt) is added to a solvent (e.g., water), a solution forms.
• A solution is a homogeneous mixture of two or more substances.
• The solvent dissolves the solute, forming a uniform mixture.
• Water's polarity allows it to dissolve many ionic compounds (e.g., salt).
• Water's positive hydrogen ends attract negative ions, while its negative oxygen ends attract positive ions, separating and dissolving the ions.
• Water can also dissolve polar molecules (e.g., sugar) through hydrogen bonding.
• Even large molecules with ionic or polar regions can dissolve in water.

Acids, Bases, and pH

• In liquid water, a tiny fraction of water molecules dissociates into hydrogen ions (H⁺) and hydroxide ions (OH⁻).
• These ions are highly reactive, and fluctuations in their concentrations can significantly impact a cell's proteins and other complex molecules.
• Some chemical compounds add H⁺ to an aqueous solution, while others remove H⁺.
• A substance that donates hydrogen ions to solutions is called an acid.
• Hydrochloric acid (HCl) in stomach gastric juice is a strong acid.
• An acidic solution has a higher concentration of H⁺ than OH⁻.
• A base reduces the hydrogen ion concentration in a solution.
• Some bases, like sodium hydroxide (NaOH), donate OH⁻ ions, combining with H⁺ to form water (H₂O), lowering the H⁺ concentration.
• Sodium hydroxide is commonly found in oven cleaners.
• Other bases increase the OH⁻ concentration by accepting H⁺ ions from the solution.
• The pH scale measures how acidic or basic a solution is.
• It ranges from 0 (most acidic) to 14 (most basic), with each pH unit representing a 10-fold change in H⁺ concentration.
• Pure water and neutral aqueous solutions have a pH of 7, with equal concentrations of H⁺ and OH⁻ ions.
• The pH inside most cells is close to 7, while human blood has a pH of about 7.4.
• Blood pH must remain within a narrow range (7.0 to 7.8) for survival.
• Biological fluids contain buffers that help maintain a constant pH by accepting or donating H⁺ ions as needed.

Explore the fundamental concepts of matter and elements in this quiz. Learn about the different states of matter, the role of chemical elements, and the formation of compounds. Understand how these elements are essential for living organisms and their unique characteristics.