Matter & Its Properties PDF

Flammability: Ability to burn (e.g., MATTER & IT’S PROPERTIES wood burns). Reactivity: Substance reacts Classification of Matter: chemically (e.g., iron rusting). 1. Physical Properties:...

Flammability: Ability to burn (e.g., MATTER & IT’S PROPERTIES wood burns). Reactivity: Substance reacts Classification of Matter: chemically (e.g., iron rusting). 1. Physical Properties: Classification of Matter: ○ Observable without changing the substance. 1. Pure Substances: ○ Examples: Height, weight, ○ Definite composition and volume, temperature, shape. properties. ○ Extensive Properties: ○ Elements: Simplest form, Depend on the amount of composed of one type of matter (e.g., mass, volume). atom. ○ Intensive Properties: ○ Compounds: Composed of Depend on the type of matter molecules made from (e.g., boiling point, density). different elements in fixed 2. Chemical Properties: proportions. ○ Observable only when the 2. Mixtures: substance's identity changes. ○ Combination of substances. ○ Examples: Flammability, ○ Homogeneous Mixtures: reactivity, combustibility. Uniform composition (e.g., solutions). Physical Properties Explored: ○ Heterogeneous Mixtures: Distinct phases or parts. Density: 3. Types of Bonds: ○ Amount of mass in a given ○ Ionic: Electron transfer volume. Formula: D = m/v. (metal + non-metal). Malleability: ○ Covalent: Electron sharing ○ Ability to be pounded into (non-metal + non-metal). thin sheets. ○ Metallic: Between metallic Ductility: elements. ○ Ability to be drawn into a wire. States or Phases of Matter: Solubility: ○ Ability to dissolve in another 1. Solids: substance (e.g., sugar in ○ Definite shape and volume. water). ○ Particles are packed tightly Thermal Conductivity: and vibrate in place. ○ Ability to transfer heat (e.g., 2. Liquids: metals conduct heat well). ○ No definite shape but have definite volume. Chemical Properties Examples: ○ Particles move around each other. 3. Gases:  ○ No definite shape or volume. COMMON SEPERATION ○ Particles move freely with TECHNIQUES high energy. 4. Plasma: Methods of Mixture Separation: ○ Electrically charged gas-like state. 1. Mechanical Separation: ○ Extremely high energy ○ Takes advantage of physical particles (e.g., lightning, neon properties like color and lights). shape. ○ Example: Recycling Changes in States of Matter: materials like plastic, paper, and metal by sorting them Adding or subtracting energy causes based on appearance. changes between solid, liquid, and 2. Magnetic Separation: gas. ○ Utilizes the physical property ○ Solid + Energy = Liquid of magnetism to separate ○ Liquid + Energy = Gas metals from non-metals. ○ Example: Separating metals Types of Changes: in a scrap yard using magnets. 1. Physical Change: ○ Turbo Beads Animation: ○ Affects physical properties, Magnetic beads are used to but no new substances separate impurities from formed. water. ○ Example: Melting butter. 3. Filtration: 2. Chemical Change: ○ Uses the physical property of ○ Produces new substances the state of matter (solid vs. with different properties. liquid). ○ Example: Rust formation on ○ A screen or filter traps solid iron, burning wood. particles while allowing liquids to pass through. Signs of a Chemical Change: ○ Examples: Filtering coffee, sieving sand, or using air 1. Odor Production (e.g., rotting food). filters. 2. Change in Temperature (exothermic 4. Decanting: or endothermic reactions). ○ Involves pouring off a liquid, 3. Change in Color (e.g., fruit ripening). leaving behind a solid or 4. Formation of Bubbles (e.g., gas another liquid based on release during chemical reactions). density differences. 5. Formation of a Precipitate (solid ○ Example: Decanting water produced from two liquids). from rice or separating a liquid from a precipitate. 5. Distillation:  ○ Separation of liquids based 10. Fractional Crystallization: on different boiling points. Involves the crystallization of ○ Example: Distillation of substances from a solution when alcohol or crude oil. their solubility limit is reached as the ○ Distillation Animation: solution cools. Demonstrates how two Examples: Growing rock candy or liquids can be separated crystallization in magma chambers. based on their different Fractional Crystallization boiling points. Animation: Shows how crystals 6. Evaporation: form at different temperatures as ○ Involves vaporizing a liquid, magma cools. leaving dissolved solids behind. ○ Example: Sea salt is obtained by evaporating ATOMS, MOLECULES, AND seawater. IONS 7. Density Separation: ○ More dense components sink, and less dense Laws of Chemical Changes: components float. 1. Law of Conservation of Mass: ○ Example: Oil and water ○ Antoine Lavoisier separation due to density discovered that the mass of differences. substances remains constant 8. Centrifuge: in a closed system during ○ Uses circular motion to help chemical reactions. denser components settle at ○ Example: the bottom faster. When 24g of ○ Examples: Separation of magnesium (Mg) blood components or DNA reacts with 71g of extraction from blood. chlorine (Cl₂), 95g of 9. Paper Chromatography: magnesium chloride ○ Separates mixtures using (MgCl₂) is produced. molecular polarity and This illustrates that attraction to different phases. the mass of reactants ○ Example: Separation of equals the mass of plant pigments or dyes on the products. paper. ○ More Examples: ○ Chromatography Water Formation: 1g Animation: Illustrates how of hydrogen (H₂) different molecular reacts with 8g of components separate based oxygen (O₂) to form on their attraction to the 9g of water (H₂O). solvent or stationary phase Iron Sulfide (paper). Formation: 5.58g of  iron (Fe) reacts with compounds like NO 3.21g of sulfur (S) to (nitrogen monoxide), produce 8.79g of NO₂ (nitrogen iron(II) sulfide (FeS). dioxide), and N₂O₄ Magnesium Oxide (dinitrogen tetroxide), Formation: 2.43g of with N magnesium burns to mass ratios of 14:16, produce 4.03g of 14:32, and 14:64 magnesium oxide respectively. (MgO), with 1.6g of oxygen reacting. 2. Law of Definite Proportion: ○ Discovered by Joseph Atomic Structure: Proust, this law states that a chemical compound always 1. Atomic Number (Z): contains the same proportion ○ Represents the number of of elements by mass, protons in an atom's nucleus. regardless of the compound's ○ Also indicates the number of source. electrons in a neutral atom. ○ Examples: ○ Example: Carbon (C) has 6 Water from any protons, so its atomic source (river, ocean, number is 6. or rain) has a 2. Mass Number (A): constant 1:8 mass ○ The sum of protons and ratio of hydrogen to neutrons in an atom’s oxygen. nucleus. Problem Example: If ○ Example: Oxygen has 8 1g of hydrogen protons and 8 neutrons, so combines with 8g of its mass number is 16. oxygen, how much 3. Isotopes: hydrogen would react ○ Atoms of the same element with 10g of oxygen? with the same number of 3. Law of Multiple Proportion: protons but different numbers ○ Proposed by John Dalton, of neutrons. this law explains how two ○ Example: Carbon has three elements can form different isotopes: Carbon-12, compounds. The masses of Carbon-13, and Carbon-14, one element that combine all with 6 protons but varying with a fixed mass of the other numbers of neutrons. element form simple 4. Neutrons: whole-number ratios. ○ Neutrons are uncharged ○ Example: particles in the nucleus. Nitrogen and oxygen ○ The number of neutrons can form multiple be found by subtracting the  atomic number from the ○ Example: The empirical mass number. formula of glucose ○ Example: (C6H12O6) CH2O Helium (He): Atomic number = 2, Mass number = 4 → 4 - 2 = Naming of Compounds: 2 neutrons. 1. Ionic Compounds: ○ Composed of cations and anions. The cation is named Ions: first, followed by the anion with an “-ide” suffix. 1. Formation of Ions: ○ Example: NaCl is named ○ Atoms become ions when sodium chloride. they lose or gain electrons. 2. Binary Compounds: ○ Cations are positively ○ Made of two elements. charged ions (loss of Prefixes like mono-, di-, tri- electrons), while anions are are used to denote the negatively charged (gain of number of atoms. electrons). ○ Example: CO2​is carbon 2. Monoatomic Ions: dioxide. ○ Ions consisting of a single 3. Acids: atom. ○ Binary acids: Composed of ○ Example: Sodium ion (Na⁺), hydrogen and one other chloride ion (Cl⁻). non-metal element. Named 3. Polyatomic Ions: using the prefix “hydro-” ○ Ions composed of two or followed by the element’s more atoms bonded together. name with an “-ic” suffix, and ○ Example: Nitrate (NO₃⁻), ending with “acid.” sulfate (SO₄²⁻). ○ Example: HCI is hydrochloric acid. Chemical Formulas: Calculating Empirical and 1. Molecular Formula: Molecular Formulas: ○ Provides the exact number of atoms of each element in a Steps: molecule. ○ Convert percentages to ○ Example: Glucose has the grams (assume 100g). molecular formula C6H12O6​. ○ Convert grams to moles 2. Empirical Formula: using molar masses. ○ Gives the simplest ○ Divide by the smallest whole-number ratio of atoms number of moles. in a compound.  ○ Round or multiply to get Chemical Equations whole numbers. ○ Example 1: A compound A chemical equation is a shorthand with 13.5% calcium, 10.8% notation to describe a chemical oxygen, and 0.675% reaction, with reactants on the left, hydrogen has the empirical products on the right, and arrows formula Ca(OH)2 indicating the direction of the Molecular Formula: reaction. ○ Derived by comparing the molecular mass to the Types of Chemical Reactions empirical formula mass. ○ Example: For glucose, if the 1. Decomposition: A single compound molecular mass is 180g/mol breaks down into two or more and the empirical formula substances mass is 30g/mol, the 2. Synthesis (Combination): Two or molecular formula is C6H12O6 more reactants form a single product 3. Single Displacement: One element STOICHIOMETRY replaces another in a compound 4. Double Displacement: Two ionic Key Concepts of Stoichiometry compounds exchange ions 5. Combustion: A hydrocarbon reacts Stoichiometry is the study of the with oxygen to form carbon dioxide quantitative relationships in and water substances and their reactions, involving chemical equations, the Balancing Chemical Equations mole, molar mass, chemical formulas, and mass relationships in Steps include writing correct equations. formulas for substances and adding coefficients to ensure equal numbers Important Definitions of each atom on both sides of the equation. Reactants: Substances consumed in a reaction. Mole Concept and Molar Mass Products: Substances formed in a reaction. Mole: A unit measuring the amount Coefficients: Numbers before the of substance, containing Avogadro’s formula of a substance, indicating number - (6.022 X 10 to the power of 23) the number of molecules or moles. Molar Mass: The mass of one mole Balanced Equation: An equation of a substance. with the same number of atoms of each element on both sides. Stoichiometric Calculations ○ Gases are compressible with low densities, as they contain Convert between grams, moles, and a lot of empty space. molecules using molar mass and ○ Gases undergo diffusion Avogadro’s number. (movement from high to low Identify the limiting reactant and concentration) and effusion calculate the amount of products (gas escaping through small formed. holes). Percent Yield: A measure of the efficiency of a reaction, calculated as Kinetic Molecular Theory (actual yield over theoretical yield) × 100%. Describes the behavior of ideal gases: Example Calculations ○ Particles have no volume and undergo elastic Limiting Reactant: Determine collisions. which reactant is used up first in a ○ Particles are in constant, reaction, thereby limiting the amount random, straight-line motion. of product formed. ○ There are no attractive or Theoretical Yield: The maximum repulsive forces between amount of product expected based particles. on stoichiometric calculations. ○ The average kinetic energy Percent Yield: Compares the actual of particles is directly yield obtained from a reaction with proportional to temperature. the theoretical yield. Real Gases vs. Ideal Gases GASES & GAS LAWS Real Gases: Key Concepts and Behavior of ○ Have their own volume and attract each other Gases (intermolecular forces). ○ Behave most ideally at low States of Matter: The state of pressures and high matter is determined by forces temperatures. (intermolecular and intramolecular) ○ Deviations occur at high and kinetic energy. Gases have high pressures and low kinetic energy, leading to particle temperatures due to motion that overcomes attractive significant intermolecular forces. forces and volume effects. Characteristics of Gases: ○ Gases expand to fill any container due to random Pressure and Measurement motion and no attraction Pressure is the force exerted per between particles. unit area. Atmospheric pressure is  exerted by gas molecules pulled (directly proportional) at toward Earth by gravity. constant volume: Units of Pressure: P1T2=P2T1 ○ 1 atm = 101.325 ○ Example: Calculating ○ kPa = 760 pressure changes in a ○ mm Hg = 760 canister as temperature ○ torr = 14.7 psi. changes. Standard Temperature and 4. Combined Gas Law: Pressure (STP): 0°C (273 K) and 1 ○ Combines Boyle’s, Charles’, atm. and Gay-Lussac’s laws: P1V1T2=P2V2T1 Temperature and the Kelvin Scale ○ Example: Finding gas volume at different conditions Kelvin Scale: Absolute temperature of pressure and temperature. scale with 0 K as absolute zero, 5. Ideal Gas Law: where particle motion stops. ○ Relates pressure, volume, Conversion formulas: temperature, and the number ○ K=°C+273K = °C + of moles of a gas: PV=nRT 273K=°C+273 where R is the gas constant ○ Always use Kelvin for gas (0.0821 L·atm/mol·K). law calculations. ○ Example: Calculating the volume occupied by moles of Gas Laws a gas at specific conditions. 1. Boyle’s Law: Applications and Example ○ Relates pressure and volume Calculations of a gas (inversely proportional) at constant Calculations involve manipulating temperature: P1V1=P2V2 equations to solve for unknown ○ Example: Calculating volume variables based on given conditions changes of a weather balloon (e.g., pressure, volume, as it rises to different temperature). altitudes. Understanding these laws allows for 2. Charles’ Law: predictions of gas behavior under ○ Relates volume and absolute varying conditions, essential for temperature (directly practical applications like scuba proportional) at constant diving tanks, weather balloons, and pressure: V1T2=V2T1 more. ○ Example: Calculating changes in a balloon's __________________________________ volume with temperature. 3. Gay-Lussac’s Law: End of reviewer, good luck hehe! - ○ Relates pressure and Trixia