CHEMISTRY: The Study of Matter

Describing Matter Chemistry: study of matter Matter: has mass and takes up space Types of matter: pure substances (sugar, salt) not pure substances (milk, fruit)

States of Matter: GASES Volume can change easily Gas particles spread apart filling all of the space available Do not have a definite shape or a definite volume

GAS BEHAVIOR Boyle’s Law: Increased pressure= decreased volume. Decreased pressure= increased volume.

GAS BEHAVIOR Charles’ Law: Increased temperature = increased volume Decreased temperature = decreased volume

IN SUMMARY…. Gases at high temperatures have high pressure and a larger volume Gases at low temperatures have low pressure and low volumes.

States of Matter: SOLIDS AND LIQUIDS Solids: definite shape, definite volume. Liquids: definite volume, but no shape solids liquids

States of Matter: SOLIDS AND LIQUIDS Like gases, solids and liquids expand when heated and shrink when cooled. What evidence do you have of this? Can you think of an exception?

Changing State deposition melting evaporating condensing freezing sublimation

Describing Matter Atoms: The most basic particles that make up matter Elements: pure substances Cannot be broken down into smaller pieces protons neutrons

Describing Matter Compounds: two or more elements chemically combined (NaCl,sodium chloride, table salt) Mixtures: two or more substances, not in a specific pattern (salad dressing, dirt). can be homogeneous (same throughout- salt water) or heterogeneous (not the same throughout- salad dressing)

Measuring Matter Weight: how gravity affects the matter in an object Mass: the amount of matter in an object, not affected by gravity (unit= gram) Volume: the amount of space matter occupies (unit = milliliter, liter) Density: how tightly matter is packed into a given space (unit = gram / ml or l) Density = Mass / Volume

Describing Matter Two Properties of Matter: Physical Properties: observed without the matter changing (state of matter, color, texture, flexibility, solubility) Chemical Properties: observed when matter changes into new substances (flammability, reactivity)

Changes in Matter Physical Changes: changes the form or appearance of matter, but does not make the matter into a new substance (ice melting) Phase changes are physical changes deposition melting evaporating condensing freezing sublimation

Changes in Matter Chemical Changes: Produces new substances (candle burning).

Density Density relates the mass of a material in a given volume. Density = Mass ÷ Volume (g/mL) Objects with a HIGH density will sink in water Objects with a LOW density will float in water

Atoms and the Periodic Table Atom: smallest particle of an element Have three, subatomic particles: protons (+), neutrons (), and electrons (-). Protons and neutrons: in the center of an atom (nucleus) surrounded by electrons in a cloud.

Scale and Size of Atoms Amazingly small Tiniest speck of dust contains 10 million billion atoms

Atomic Number Every atom of an element has the same number of protons. Unique number of protons = atomic number. Atomic number identifies an element. Carbon (C) = 6, Oxygen (O) = 8, Iron (Fe) = 26. If an atom has to be stable, how can you figure out the number of electrons?

Organization of the Periodic Table Properties of an element can be predicted from its location on the PT Periods: Horizontal rows on the PT (elements lose reactivity as you move across periodic table, until you reach the gases) Groups: Vertical columns on the PT (elements in groups have similar properties)

Reading an Element’s Square

Metals Physical Properties: shiny, malleable (bendable), ductile (stretchy), conductive (transfers heat or electricity) Chemical Properties: react by losing electrons to other atoms (the fewer electrons to lose or gain, the more reactive an element is)

Metals Some are extremely reactive (sodium) others are not reactive at all (gold, platinum) Some have moderate reactivity (iron- corrosive) Reactivity decreases as you move from left to right across the periodic table

Alkali Metals Group 1: react with other elements by losing one electron. -So reactive, never found as uncombined elements in nature -Two most important: Sodium and potassium

Alkaline Earth Metals Group 2 React by losing two electrons Two most common are magnesium and calcium Each is fairly hard, gray-white, and a good conductor

Transition Metals Metals in Groups 3-12 Most of familiar metals (iron, platinum, copper, nickel, silver, gold) Most are hard and shiny, good conductors Less reactive as you move to the right

Misc. Metals Some elements in groups 13 and 14 (aluminum, tin, lead) Not very reactive Lanthanides (top row in rows below PT): soft, malleable, shiny, highly conductive). Actinides (row below Lanthanides): Only Actinium (Ac), Thorium (Th), Protactinium (Pa) and Uranium (U) occur naturally

Nonmetals Nonmetal: lacks most of the properties of a metal (not conductive, reactive, dull, brittle) Carbon Family (group 14): Elements can gain, lose, or share four electrons Very versatile

Nonmetals Nitrogen Family (group 15): Gain 3 electrons Occur in nature in the form of diatomic molecules (Di = two, atomic = of atoms). Ex: P2,N2,

Nonmetals Oxygen Family (group 16): Gains or shares 2 electrons Because oxygen is highly reactive, it can combine with most other elements in nature Oxygen you breathe is diatomic O2, or can be triatomic, O3, ozone.

Nonmetals Halogen Family (group 17): Halogen means salt-forming (bonds to metals to form salts) Gains 1 electron Very reactive, dangerous (but can bond to other elements to be useful)

Nonmetals Noble Gases (group 18): Stable Exist in the Earth’s atmosphere in small amounts Don’t bond to other elements (have enough electrons)

Hydrogen Simplest element with the smallest atoms Unique properties, cannot be grouped with other gases Makes up 1% of Earth’s crust Rarely found on Earth as pure element Mostly found as H20 (water)

Metalloids Characteristics of metals and non-metals Most common is Silicon (Si) Most useful property is their varying ability to conduct electricity Semiconductors (used in computer chips) are made of metalloids

Isotopes Isotopes: identified by the mass number of an element (the number of protons plus neutrons) Carbon is most commonly C-12, but can be C-13 and C-14. What’s the difference?

Isotopes and Mass Number Atoms of an element always have the same number of protons, but can have different numbers of neutrons. Atoms of an element with different numbers of neutrons: isotopes

Phase Changes and Energy The phase a substance depends on the kinetic energy of the molecules Substances change from phase to phase in response to the input or output of energy Phase changes reflect a change in energy or a change in the behavior of the atoms or molecules, not a change on the atoms or molecules themselves. Ex: ice melting, water condensing

Phase Changes and Energy Matter can change from one state to another when thermal energy is absorbed or released.

Law of Conservation of Energy During any chemical or physical change, energy cannot be created or destroyed. Energy can change forms and can be transferred from one object to another, but the overall amount must be conserved.

Temperature A measure of the average kinetic energy of individual particles of matter Measured using a thermometer (heat increases the kinetic energy of the liquid inside, takes up more space).

Thermal Energy and Heat The total energy of all of the particles of an object is called Thermal Energy. The more particles an object has at a given temp, the greater the thermal energy. The transfer of thermal energy from matter at a higher temperature to matter at a lower temperature is called Heat. Why does ice melt in your hands?

Heat Transfer Heat is transferred in 3 ways: Conduction: Heat is transferred without the movement of the matter (metal handle of a pot getting hot). Convection: Heat is transferred by the movement of currents within a fluid (plate tectonics). Radiation: The transfer of energy by electromagnetic waves. Radiation does not require matter to transfer energy.

Examples of Heat Transfer

Direction of Heat Transfer Heat will flow from the warmer object to the colder object Heat transfer occurs in only one direction In summary, heat always flows from higher energy levels to lower energy levels.

Phase Changes and Energy Energy transferred TO the substance deposition melting evaporating freezing condensing sublimation Energy transferred FROM the substance

Calories Heat is measured in Calories. 1 Calorie is the amount of heat needed to raise the temperature of 1gram of water 1 degree Celsius. Equilibrium: when there is no net energy transfer within a system

Heat of Fusion Water molecules in ice are held together by forces called “bonds.” Energy is required to break the bonds when changing solid water (ice) to liquid water. The energy needed to change solid water to its liquid form is called Heat of Fusion. The Heat of Fusion for water = 80c/g (calories per gram)

Observing Chemical Change Matter has chemical and physical properties and undergoes chemical changes and physical changes.

Chemical Changes Occur When… Chemical changes occur when bonds break to form new bonds Chemical changes occur through chemical reactions

Evidence for Chemical Reactions Two main changes occur during a chemical reaction: 1) Formation of new substances 2) Changes in energy

Evidence for Formation of New Substances Color Change Precipitate Formed (solid) Gas Produced

Evidence for Changes in Energy Heat Produced (reactions can be exothermic or endothermic)

Matter and Energy When a reaction causes energy to be released (often as heat), the reaction is called exothermic. When energy is taken in, a reaction is called endothermic (ice melting). Endothermic Energy taken in, resulting temp is higher Exothermic Energy released, resulting temp is lower

Describing Chemical Reactions Chemical Equations: A short, easy way to show a chemical reaction, using symbols instead of words. Structure of Chemical Equations: All chemical equations have a common structure. Substances at the beginning are called “reactants” and substances at the end are called “products”

Chemical Equations reactant + reactant  product + product You read the arrow () as “yields.”

Law of Conservation of Mass Chemical Equations must be balanced because of the Law of Conservation of Mass (matter can neither be created nor destroyed)

Steps for Balancing Equations 1) Write the equation H2 + O2  H2O reactants products Count the atoms H2 = 2 hydrogen atoms (reactant) O2 = 2 oxygen atoms (reactant) H2 = 2 hydrogen atoms (product) O = 1 oxygen atom (product)

Steps for Balancing Equations Use coefficients to balance atoms (a coefficient is a number placed in front of a chemical formula to balance products and reactants) H2 + O2  2H2O **Coefficients can be placed on the reactant side or the product side of an equation 4) Look back and check

Checking your Equations

Chemical Reactions and Bonding Bonds are broken and formed during chemical reactions. If an atom loses or gains an electron it is called an ion (an element with a + or – charge). Bonds that form between ions are called ionic bonds (usually between metals and nonmetals).

Ionic Bonding

Chemical Reactions and Bonding If elements bond by sharing electrons, covalent bonds form. Covalent bonds usually form between nonmetals.

Covalent Bonds

Solutions Solutions: a well-mixed mixture. Solvent: What does the dissolving (the larger part of a solution). Solute: What is dissolved (the smaller part of a solution). Liquid water solution

Solutions The effects of solutes on solutions: Lower freezing points, Higher boiling points. What affects solubility? Temperature, pressure, type of solute and solvent

Concentration of a Solution Saturated Solution: A solution that contains as much solute dissolved as possible. (No more can be dissolved) Unsaturated Solution: A solution where more solute can still be dissolved.

Acids and Bases Properties of Acids: tastes sour, reacts with metals and carbonates, and turns litmus or pH paper red. Uses of Acids: vitamins, batteries (sulfuric acid), fertilizers

Acids and Bases Properties of Bases: taste bitter, slippery, turns indicator paper blue. Opposite of acids: Do NOT react with metals and carbonates Uses of Bases: Cleaning Products

Acids and Bases The pH scale used to determine the strength of an acid or a base. The pH scale measures the concentration of hydrogen ions in a solution. A low pH: very acidic. A neutral solution (water) pH = 7 A high pH: very basic.

Acid-Base Reactions When an acid and base react, the product is neither acidic or basic. Neutralization occurs when an acid and base react (pH becomes closer to neutral) Salts form as a product of an acid base reaction when the positive ion in a base bonds with a negative ion of an acid. Ex: HCl + NaOH  H2O + (Na+ + Cl-)