Double Displacement Precipitation - the exchange of 2 ionic compounds they precipitate AB + CD = AD + BC - reactions that form solid are precipitate and the aqueous solutions are the precipitation reactions - The only ionic compounds that can precipitate are insoluble compounds. Atoms -All matter are made of atoms - Proton – has positive charge -Neutron – no charge -Electron – negative charge, has no real mass -Atomic Number – no. of protons in an atom -Neutral atom – no. of protons = no. of electrons -Atomic Mass- no. neutrons and protons -No. of protons in an atom never changes but neutrons can change -Isotopes – atoms of the same element but with a different no. of neutrons - Ions –atoms of the same element but with a different no. of electrons *If an atom loses electrons, it becomes positively charged * If an atom gains electrons, it becomes negatively charged Atomic Theory 1. Early Theory - Aristotle – all matter are made of 4 elements fire, air, water, earth - Democritus – all matter are made of small units called atomos 2. Modern Theory - Dalton (1800 England) * All elements are made of atoms and they are indestructible solid spheres. They can’t be created or destroyed THE LAW OF CONSERVATION OF MATTER *Atoms of the same element are exactly alike. *Atoms of different elements are different compounds formed by joining of atoms of two or more elements - Thomson (1897-England) *Discovered the electrons -Rutherford (1897- England) *1908, discovered the Nucleus -Bohr(1913 Denmark) *Electron are in orbits and has certain energy levels around the nucleus Electron Cloud Model - Electrons travel in regions called electron clouds - You can’t predict the exact location of electrons Energy Levels -The energy of an electron is based on its location around the nucleus -Electrons that are closer to the nucleus has less energy Level 1 – 2 electrons Level 2 – 8 electrons Level 3 – 8 electrons Level 4 -18 electrons Level 5- 32 electrons Heat -transfer of energy due to temperature changes Exothermic * when the system gives off heat usually the system is hotter than the surroundings Endothermic *when the system is cooler than the surroundings Calorimetry -the measure of heat flow -Calorimeter *used to measure Calorimetry Specific Heat Capacity- amount of heat needed to raise the temperature of 1g of substance by 1 degree Celsius Energy changes heat of fusion (ΔHfus): melting/freezing heat of vaporization (ΔHvap): boiling/condensing Equation of Heat q = m C ΔT (used within a given state of matter) q = m ΔH (used during phase change) where q = heat (J) C = specific heat (J/g°C) ΔH = heat of vap/fus (J/g) m = mass of substance (g) ΔT = temperature changes (°C) Organic Compounds and the Atomic Properties of Carbon - Carbon forms covalent bonds in all its elemental forms and compounds. - Carbon can CATENATION (the ability to bond to itself to form stable chain/ring/branched compounds.) - A Carbon Carbon bond (C-C) is much stronger than Silicon Silicon bond (Si-Si) - Heteroatoms (atoms other than a C and H that is part of a organic compound) - Carbon atoms can form a maximum of 4 bonds Hyrdocarbons - compounds that contain only C and H - 6 C atoms, no multiple bonds, no rings -4 C atoms, one double bond, no rings -4 C atoms, no multiple bonds, one ring. Alkanes - Hydrocarbons that contain only a single bond, suffix -ane, CnH2n+2 Alkenes - hydrocarbon that contains atleast one C=C bond, suffix alkenes -ene, CnH2n Naming Organic Compounds - Prefix identifies any groups attached to the main chain - Root identifies to the no. of C atoms - Suffix indicates the type of organic compound meth- Root no. C atoms 1 eth- 2 prop- 3 but- 4 pent- 5 hex- 6 hept- 7 oct- 8 non- 9 dec- 10 - Constitutional or Structural isomers. same molecular formula but different arrangement bonded atoms. - Stereoisomers. same arrangement of atoms but different orientations of groups in space. - Optical Isomers . mirrored images of each other that can’t be superimposed. - A compound that rotates the plane of light clockwise is called dextrorotatory, while a compound that rotates the plane of light counterclockwise is called levorotatory. - Geometric Isomers, alkene which differ in the orientation of the groups attached to the double bond. Alkynes - hydrocarbons that contains at least one CΞC triple bond, suffix –yne. , CnH2n-2 -A saturated solution is a solution that contains the maximum amount of solute that is capable of being dissolved. - the electron rich –CΞC– group Type of Organic Reactions - addition reaction occurs when an unsaturated reactant becomes a saturated product - elimination reaction occurs when a saturated reactant becomes an unsaturated product, The groups typically eliminated are H and a halogen atom or H and an –OH group. -- substitution reaction occurs when an atom or group from an reagent substitutes for one attached to a carbon in the organic reagent, The C atom at which substitution may be saturated or unsaturated, and X and Y can be many different atoms. Function Groups Alcohols - carbon bonded to –OH group, ends in –ol, high melting and boiling points Reaction of Alcohols - dehydration(elimination) -oxidation(elimination) Aldehydes and Ketones Haloalkanes - carbon bonded to a halogen atom Amines - carbon bonded to nitrogen atom, ending –amine. Amides - O=C–N, functional group, can be hydrolyzed to give a carboxylic acid and an amine. -both contain the carbonyl group, C=O -Aldehydes , ends in –al -Ketones, ends in –one Organometallic - metal atom covalently bonded to C Carboxylic Acids - functional group –COOH, ending –oic acid Esters - formed by a reaction of alcohol and carboxylic acid -common in lipids Saponification - process of converting fat, oil/lipid into soap by using a base solution Polymers - long continuous molecules that repeats with each other - Addition Polymers/Chain-growth polymers *monomers that go through addition reaction with each other (most addition polymers contain an alkene group) - Condensation polymer *monomers that go through dehydration-condensation with each other (condensation polymers have two functional groups, and each monomer can link to two others.) Equilibrium - All reactions are reversible and under certain conditions can reach a state of equilibrium. -Equilibrium- forward rate = reversal rate - Chemical Equilibrium, dynamic state, the reactions can still occur while being in a equilibrium but because they occur at the same rate, no net change is observed on the macroscopic level. - K =relationship of products and reactants of a reaction at a equilibrium. - If the value of K - small, reaction yields little product before reaching equilibrium, reactants Reaction Quotient Q -large, reaction reaches equilibrium with little reactant remaining, products - aA + bB = cC + dD -Qc = [𝐶]𝑐 [𝐷]𝑑 [𝐴]𝑎 [𝐵]𝑏 -Q gives the ratio of product concentrations to reactant concentrations at any point in a reaction. - For a particular system and temperature, the same equilibrium state is attained regardless of starting concentrations. The value of Q indicates how close the reaction is to equilibrium, and in which direction it must proceed to reach equilibrium. Forms of K and Q -For an overall reaction that is the sum of two more individual reactions: Qoverall = Q1 x Q2 x Q3 x ….. Koverall = K1 x K2 x K3 x …… Heterogeneous equilibrium - reactants and/or product in different phases -A pure solid or liquid always has the same “concentration”, i.e., the same number of moles per liter of solid or liquid. -The expressions for Q and K include only species whose concentrations change as the reaction approaches equilibrium. Pure solids and liquids are omitted from the expression for Q or K.. - If Q < K, the reactants must decrease and the products increase; the reaction proceeds toward the products; reactants → products until equilibrium is reached. - If Q > K, the reactants must increase and the products decrease; the reaction proceeds toward the reactants; reactants ← products until equilibrium is reached. - If Q = K, the system is at equilibrium and no further net change takes place. Le Châtelier’s Principle - chemical system at equilibrium is disturbed, it reattains equilibrium by undergoing a net reaction that reduces the effect of the disturbance - The system responds to a disturbance by a shift in the equilibrium position - Shift to the left is a net reaction of product to reactant, shift to the right is a net reaction from reactant to product Electrochemistry - unites chemistry and electricity, study of the transfer of electrons - Voltage/Charge/Current - electric potential difference is called voltage. Measurement of voltage is volts - Coulomb, measurement of electric charge -Current, the flow of charge per unit of time, Ampere measurement of current, 1C/s -Energy (J)= potential (V) x Charge(C) Cathode /Reduction and Anode/Oxidation - When the electronic conductor gives up electrons to the ionic conductor then the reaction is called a cathodic reaction. - Electrode that produces electrons are called anodic reaction. Electrons flow in the opposite direction to the positive side - Oxidation, the loss of electrons - Reduction, the gain of electrons -Anode, the electrode where electricity moves into, the positive side -Cathode, the electrode where electricity is given out or flows out of, the negative side Electrochemical Cell -electrons produced by an anode, by cathodic reaction. -anode must be paired with a cathode, combine in a electrochemical cell -created by placing metallic electrodes into an electrolyte where a chemical reaction either uses or generates an electric current - Electrochemical cells which generate an electric current ( in which the 2 electrode reactions occur spontaneously) are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells. Electrolytic Cell -chemical cells used to do electrolysis, containing electrolyte, usually a solution of water or other solvents that can dissolve ions. Nuclear Chemistry - Nucleus is composed of nucleons, protons and neutrons - atomic mass number (A ; total no. nucleons) - atomic number (Z ; number of protons) -ZAE - Carbon-12 has a mass of 12 u, 6 protons and 6 neutrons Subatomic Particles - (u) atomic mass unit, 1u = 1/12th the mass of a carbon-12 atom - Proton and Neutron’ s mass > 1u - electron, 9.11 x 10-31 kg, 5.485 x 10-4 u - proton, 1.673 x 10-27 kg, 1.0073 u - neutron, 1.675 x 10-27 kg, 1.0087 u - The mass of a nucleus is slightly less than the mass of the individual nucleons - Missing mass of the nucleus is called mass defect, Δm = mass of the protons and neutrons – mass of the nucleus - Energy and mass can be interconverted - E = mc2 - When protons & neutrons are packed together to form a nucleus, some of the mass is converted to energy and released - This amount of mass is equal to the force of attraction holding the nucleons together -The total energy required to break up a nucleus into its constituent protons & neutrons - The number of neutrons in a nucleus can vary - too much or too few neutrons, can cause instability - Stable nuclei doesn’t decay spontaneously - Unstable nuclei have a certain probability to decay Radioactivity - spontaneous decomposition of an unstable nucleus into a more stable nucleus by releasing energy Electromagnetic Radiation - form of energy that can pass through empty space - electromagnetic radiation can be a particle and a wave - the shorter the wavelength, the more energy it possesses - Alpha Decay (increases N:Z ratio) - Beta Decay (decreases N:Z ratio) - Gamma Decay - Alpha Particles, relatively heavy and double charged, lose energy quickly in matter - Beta Particles, much smaller and singly charged, interact slowly in matter - Gamma rays and X-ray, high energy, more lengthy interaction with matter Hazards of Radiation Types - Alpha Emissions, easily shielded, considered hazardous is alpha emitting material is ingested or inhaled - Beta Emissions, shielded by thin layers of material, considered hazardous is beta emitting material is ingested or inhaled - Gamma Emissions - need dense material for shielding, considered hazardous when external to the body Radioactivity Decay Rates/Half Life (3 valve pls) - The relative stability of nuclei can be expressed in terms of the time required for the half of the sample to decay - time it takes for 1g to decay to .5g -Co-60 5 yr -Cu-64 13 h -U-238 4.51 x 109 yr -U-235 7.1 x 108 yr -The level of radioactivity of an isotope is inversely proportional to its half-life, the shorter the half life , more unstable the nucleus -The half-life of a radionuclide is constant -Rate of disintegration is independent of temperature or the number of radioactive nuclei present Radon-222 Parent # of Decay Steps Final Product of Series Uranium-238 14 Lead-206 Thorium-232 10 Lead-208 Uranium-235 11 Lead-207 Plutonium-241 13 Bismuth-209 Radioisotope Trends and the Periodic Table - Trends, more than 20 properties change in predictable way based location of elements Properties such as -Density - melting point/boiling point -atomic radius - ionization energy -electronegativity Atomic Radius - Atomic Radius, defined as ½ distance between neighboring nuclei in molecule or crystal - Affected by, no. of energy levels, proton pulling power Proton Pulling Power -“measured” by Proton Pulling Power by determining the Effective nuclear charge - charge felt by valence electrons - Nuclear charge- no. of inner shell electrons Ionization Energy - amount of energy required to remove a valence electron from an atom in gas phase Electronegativity - ability of atom to attract electrons in bond - noble gases tend not to form bonds , so they don’t have electronegativity values Reactivity of Non-Metals -judge reactivity of non-metals by how easily they gain electrons Reactivity of Metals - judge reactivity of metals by how easily they give up electrons How do you know if an atom gains or loses electrons? -Think back to the Lewis structures of ions -Atoms form ions to get a valence of 8 (or 2 for H) -Metals tend to have 1, 2, or 3 valence electrons , it’s easier to lose them -Nonmetals tend to have 5, 6, or 7 valence electrons, it’s easier to add some -Noble gases already have 8 so they don’t form ions very easily Positive ions (cations) - formed by loss of electrons, always smaller than parent atom Negative ions or anions - formed by gain of electrons, always larger than parent atom Allotropes - Different forms of element in same phase, different structures and properties - O2 and O3 - both gas phase, O2 (oxygen) - necessary for life, O3 (ozone) - toxic to life - Graphite , diamond , both carbon in solid form Periodic Table - 1st theory about the elements was of Aristotle’s theory that elements are; air, earth, water, fire - Scientist have discovered 90 naturally occuring elements, and created about 28 others -The elements, alone or in combinations, make up our bodies, our world, our sun, and in fact, the entire universe. - The most abundant element in the earth’s crust is oxygen, atmosphere is nitrogen -atomic number -symbol -atomic mass -no.valence electrons -state of matter at room temperature Atomic Number - no. of protons, no two elements have the same no. of protons Atomic Mass -the weight of atom, the no. of protons and neutrons Atomic Mass and Isotopes - Some atoms have more or less neutrons than protons. These are called isotopes Atomic Mass Unit - unit of measurement for an atom -1 AMU = mass of 1 proton - 1g = 6 X 1023 AMUss Valence Electrons - electrons in the outer energy level of the atom Properties of metals - metals are good conductor of heat and electricity - metals are shiny - metals are malleable - metals are ductile - metal + water = corrosion Properties of Non-metals - poor conductors of heat and electricity - not ductile or malleable - brittle and break easily, dull -most non-metals are gases Properties of Metalliods -properties of both metal and non-metals - can be shiny or dull - conduct heat and electricity better than non-metals but not as well as metals. -ductile and malleable Families - columns of elements are called groups or families -similar but not identical properties - all elements in a family have the same no. of valence electrons Periods -horizontal row of elements are called periods - not alike in properties Hydrogen (H) - gas at room temperature, one proton one electrons - only needs 2 electrons to fill its valence shell Alkali Metals -most reactive metals -reacts violently with water - never found as free elements in nature, always bonded with an another element Alkaline Earth Metals - never founded uncombined in nature - 2 valence electrons - includes magnesium, and calcium, among others Trasition Metals - metals such as copper, tin zinc, nickel, gold, silver, and iron -good conductors of heat and electricity -bright colored, used often in color paints -can form with oxygen to form oxides Transition Elements - similar properties to transition metals Boron Family - named after the element Boron -has 3 valence electrons -includes boron(a metalloid) and aluminum Carbon Family - has 4 valence electrons -includes carbon(non-metal), metalloids, and metals -Carbon (basis of life) Nitrogen Family -Nitrogen (78% of our atmosphere) -includes nitrogen, metalloids, metals - has 5 valence electrons, shares electrons when they bond Oxygen Family - has 6 valence electrons - Oxygen, most abundant element in the earth’s crust, extremely active and combines with almost all elements Halogen Family -elements includes such as; fluorine, chlorine, bromine, iodine, astatine -has 7 valence electrons, needs to gain 1 electron to fill their outermost energy level - reacts with alkali metals to form salts Noble Gases - colorless gases that are extremely unreactive - their outermost energy level is full -they don’t combine with other gases, they are called inert. - includes helium, neon, argon, xenon, krypton, radon Rare Earth Elements - 30 rare earth elements composed of the lanthanide series - most of the actinide are called trans-uranium, synthetic or man-made Mendeleev -1869, Dmitri Ivanovitch Medeléev, created the 1st accepted vers. of the periodic table Matter - composed of atoms bonded together in groups called molecules - Substance made from 1 type of atom are called pure substances - Substance made from 2 or more types of atoms are called compounds. - Compounds that are combined physically but not chemically, are called mixtures -Mixtures can be separated by physical means. - Compound can only be separated by chemical means. Kinetics - the rate at which chemical process occurs Reaction Speed- measured by the change in concentration with time Factors that can affect the rate of reations - reactant concentration - temperature - action of catalysts - surface area - pressure of the gaseous reactants and products Reaction Rates- measure the rate Rate Laws- rate depends on the amount of reactants Integrated Rate Laws – calculate the amount left of the time to reach a given amount Half-Life – how long it takes to react 50% of the reactants Arrhenius Equation- rate constant changes with temperature Mechanisms- link between rate and molecular scale processes For the reaction A B there are two ways of measuring rate: - the speed at which the reactants disappear - the speed at which the products appear Reversible Reactions- as products accumulates, they can begin to turn back into reactants Instantaneous rate =∆[A]/∆t limits to d[A]/dt Most Common Units… Rate = M/s Where Molarity (M) =moles/liter Reaction rate is the change in the concentration of a reactant or a product with time (M/s) Factors that affect the Reaction Rate Constant 1. Temperature - at high temperature, reactant molecules have more kinetic energy, moves faster, speeding or slowing up the reaction *Collision Theory, when 2 chemicals react, their molecules have to collide with each other for the reaction to take place *Kinetic Theory, increasing temperature means the molecules move faster 2. Concentrations of the reactants -As the concentrations of reactants increases so does the likelihood that reactant molecules will collide, 3. Catalysts - Speed up reactions by lowering activation energy 4. Surface area of a solid reactant -more area for reactants for contact 5. Pressure of gaseous reactants or products - Increased number of collisions Oxidation and Reduction - Oxidation-Reduction Reaction- process of transferring electrons between reactants -Oxidation- process of losing electrons or gain of oxygen -Reduction- process of gaining electrons or loss of oxygen -Reducing Agent- substance that donates electrons, it is getting oxidized -Oxidizing Agent-substance that accepts electrons, it is getting reduced Entropy and Thermodynamics -Thermodynamics- the relationships between heat and other forms of energy. - Spontaneous Reactions- exothermic reaction Entropy, S -refers to the disarrangement or randomness of atoms and molecules -PROBABILITY suggests that a spontaneous reaction will result in the dispersal. Directionality of Reactions - Matter Dispersal - Energy Dispersal -S (gases) > S (liquids) > S (solids) -Entropy increases with temperature and molecular complexity -Increases when a pure liquid/solid dissolves in a solvent Entropy Changes for Phase Changes ∆S = q/T where q = heat transferred in phase change, T = temperature (Kelvin) The total entropy of a reaction ∆So = ∑ 𝑺o (products) - ∑ 𝑺o (reactants) 2nd Law of Thermodynamics -entropy as a physical property of a thermodynamic system - ∆Suniverse = ∆Ssystem + ∆Ssurroundings -∆Suniverse > 0 for spontaneous process Reaction Types - Combination Reactions, A + B = AB - Decomposition Reactions, AB A + B - Combustion Reactions, Most often involve hydrocarbons reacting with oxygen in the air to produce CO2 and H2O Formula Weight -Sum of the atomic weights for the atoms in a chemical formula, generally reported for ionic compounds. Molecular Weight - Sum of the atomic weights of the atoms in a molecule Percent Composition -One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation Mole - 6.02x1023 of 12 C weigh 12 g - 1 mole = 6.022141x1023 things Percent Yield - A comparison of the amount actually obtained to the amount it was possible to make Intermolecular Forces - (inter = between) between molecules and the temperature (kinetic energy) of molecules. Gases: The average kinetic energy of the gas molecules is much larger than the average energy of the attractions between them Liquids: the intermolecular attractive forces are strong enough to hold the molecules close together, but without much order Solids: the intermolecular attractive forces are strong enough to lock molecules in place (high order). Types of intermolecular forces (between neutral molecules): - Hydrogen bonding: cases of very strong dipole-dipole interaction -London dispersion forces: (instantaneous dipole moment)( also referred to as van der Waal’s forces) polarizability: the ease with which an atom or molecule can be distorted to have an instantaneous dipole. “squashiness” Other types of forces holding solids together Ionic: “charged ions stuck together by their charges” , There are no individual molecules here.
