Lab Safety Being in the lab is a privilege that can be taken away. Students and teachers should conduct themselves in a responsible manner to ensure the safety of everyone General Rules: Conduct yourself in a responsible manner and check about medical conditions prior to working in the lab Follow all written and verbal instructions. If confused, ask the instructor before proceeding Never work alone- an instructor must always be present; experiments must be monitored, unauthorized experiments are prohibited Do NOT eat, drink, or chew gum in the lab Do NOT smell things directly- always waft Be prepared, no horseplay, keep areas clean Know where things are located- fire extinguisher, exites, fire alarm, etc. Work in well ventilated areas, use fume hood if necessary Be alert, proceed with caution, know what to do during an emergency Dispose of chemicals properly, read labels and equipment carefully Keep hands away from the face, mouth, eyes, nose, and body- always wash hands after experiments Be careful with sharp objects- hever catch falling instruments, cut away from body, grasp by the handle   Clothing: Wear safety goggles around: chemicals, heat, glassware No baggy or loose clothing, open toed shoes, dangling jewelry Long Hair must be tied back   Handling Chemicals and Injuries: Do not touch, smell, or taste anything unless instructed to Always check labels and dispose properly Handle chemicals with extreme care- do not drop or spill; hold securely Do not touch broken glass without gloves, use tongs or mitts for hot objects Never leave a burner unattended In case of accidents or injury, report to the instructor

The Scientific Method The Scientific Method- A systematic process for studying nature that involves observations, hypothesis, and experimentation Make an Observation State the problem Collect the data Formulate a Hypothesis Perform Experiments Types of Observations Qualitative- non-numerical; they ask "what?" Quantitative- numerical; they ask "how much?" Quantitative observations can also be called measurements   What testable questions can you come up with?

Law Vs. Theory Law Description of observed phenomena- patterns Tells what happens Theory Well-supported explanation of observations Repeatedly tested and confirmed Attmepts to explain why it happens

Accuracy Vs. Precision Accuracy- closeness of a measured value to a standard or known value Precision- the closeness of two or more measurements to each other

Significant Figures The significant figures in a measurement include all the digits that can be known precisely plus a last digit that must be estimated Rules: Every non-zero digit in a recorded measurement is significant Zero Rules: Leading Zeros- zeros that precede all the non-zero digits; NEVER count as sig-figs eg. 0.00832-- 3 sig-figs eg. 0.23-- 2 sig-figs Captive Zeros- zeros between non-zero digits; are counted as sig-figs eg. 7003-- 4 sig-figs eg. 409.076-- 6 sig-figs Trailing Zeros- zeros at the right end of a number; only count as sig-figs when there is a decimal point eg. 456200-- 4 sig-figs eg. 50938.00-- 7 sig-figs Exact Numbers Obtained by counting, not by measuring devices Can be assumed to have an unlimited amount of sig-figs Can come from definitions eg. 1 in.=2.54 cm Multiplying and Dividing sig-figs The answer must contain no more sig-figs than the measurement with the least sig=figs The position of the decimal point is irrelevant eg. 5.89 x 4.5 = 27 Adding and Subtracting sig-figs The answer must contain the same number of sig-figs to the right of the decimal point as the measurement with the fewest sig-figs to the right of the decimal point eg. 61.2 = 9.621 = 70.2   Scientific Notation Chemists use scientific notation is used for very large or small numbers eg. the mass of a hydrogen atom is 0.00000000000000000000000167 grams eg. 2.0 grams of hydrogen is composed of 602000000000000000000000 hydrogen molecules   Numbers are written as the product of two numbers: A coefficient A power of 10 with an exponent The exponent tells you how many times you multiply a number by 10 Numbers greater than 1 have a positive exponent (# of places the decimal moves left) Numbers less than 1 have a negative exponent (# of places the decimal moves right) eg. The number 23000 is written in exponential form as 2.3 x 10^4 eg. 0.000051 is written in exponential form as 5.1 x 10^-5 Numbers between 1 and 10 do not need scientific notation  eg. 9 = 9 x 10⁰ 1.2 = 1.2 x 10⁰ 7.562580 = 7.562580 x 10⁰ All have an exponent of zero (10⁰ = 1)

Matter Matter is everything that has substance and takes up space, or that has mass or volume Mass- the measure of the quantity of matter in an object Usually measured in grams (g) or kilograms (kg) 1 kg = 1000 g Volume- the measure of the quantity of matter in an object ​​​​​​​Measured in mL or cm^3 1000 mL = 1 L 1 mL = 1 cm^3​​​​​​​

Displacement Displacement is a method used to calculate the volume or an irregular object Steps: Fill the graduated cylinder with enough water to cover the object Note the starting volume using the correct number of sig-figs Slowly drop in the sample- without splashing the water Record the new volume Subtract the initial volume from the final volume to determine the volume of the sample, using the correct number of sig-figs

Density Density- the mass of a substance per unit volume An object's density is all about its mass and volume. A rock is heavy for its size. It has a high mass relative to its volume. It has a fairly high density. On the other hand, a wooden block is light for its size. It has a low mass relative to its volume. It has a fairly low density. Equation: d = m / v d- density m- mass v- volume units: g/mL or g/cm^3 REMEMBER YOUR SIG-FIGS

SI Units Le Systeme International D'Unites, abbreviated SI are object or phenomena that are of constant value, easy to preserve, and reproduce, and practical size SI has 7 base units (mass, length, time, temperature, amount of substance, electrical current, and luminous intensity)    Quantity                                                      Unit Name                                                    Unit Abbreviation Length                                                        Meter                                                                  m Mass                                                           Kilogram                                                             kg Time                                                            Second                                                                s Temperature                                             Kelvin                                                                   K Amount of Substance                              mole                                                                    mol   We also have volume... SI unit of is cubic meters, m³ Conversion: 1 cm³ = 1 mL we normally use mL or cm³

Dimensional Analysis Begin with the end in mind- what are you solving for? List your given Determine the conversion factor(s) necessary from the SI units. You may have more than one conversion factor Complete your conversion(s)    You will use this method through the entire course!

The Periodic Table Mendeleev organized his periodic table based on the properties of the elements, specifically, reactivity and atomic mass. Elements in each section of the periodic table have similar properties. Reactivity describes whether an element will chemically combine with other common substances and also describes the speed of the reaction. Mendeleev's arrangement of the elements helped predict the existence of undiscovered elements. Basic Patterns: Elements in the same group have similar properties (same number of valence electrons, etc.) Elements in the period have the same number of electron shells Atomic size increases going down Metals are on the left, non-metals on the right, other methal, metalloids, and transition metals are in the middle, halogens and noble gases to the very left, actinide and lanthanide series are listed at the bottom The number of protons corresponds to the atomic number Electronegativity and Ionization energy increase going up and right Reactivity of metals increases going down, reactivity of non-metals increases going up Group- a vertical column in the periodic table Period- horizontal rows on the periodic table Noble Gases- elements in group 8A; are not reactive (noble gases have no reaction)   Element Squares: Contain- atomic number, symbol, name, and atomic mass About the atom Diameter of the nucleus and electron cloud are measured in femtometers The nucleus is dense- its mass is made up of protons and nuetrons The electron cloud is mostly open space Electrons are identical, but have different energy levels Neutral atoms- # of protons = # of neutrons amu- atomic mass unit # of protons + # of neutrons = mass of atom in amu

Isotopes An Isotope is an atom with a different number of neutrons than the number of protons Symbols:  11/5 B (Boron-11) top number (11)- mass number bottom number (5)- atomic number   Example; Carbon will always have 6 protons, but the number of neutrons can change Carbon-12 (6 protons, 6 neutrons) Carbon-13 (6 protons, 7 neutrons) Carbon-14 (6 protons, 8 neutrons) Average atomic mass is the weighted average of the mass numbers of the isotopes of an element. This is what is listed on the periodic table.    eg. Carbon’s average atomic mass is 12.01

Nuclear Reactions Nuclear Reactions are a change in the nucleus and can change the atom into another. They often occur when unstable nuclei want to be stable. Nuclei undergo spontaneous changes in their number of protons and neutrons until they are stable. Radioactive Isotopes are unstable and go through radioactive decay, or ejecting or emitting pieces from the atom to achieve stability. Some types of radioactive decay are: Alpha decay- ejection of an alpha particle from the nucleus Beta decay- ejection of a beta particle from the nucleus Alpha Decay: Alpha particle (α) consists of two protons and two neutrons It is the same as the nucleus of a helium atom 238/92 U ------> 234/90 Th + 4/2 He (α) + (γ)  Beta Decay: A beta particle (β) is an electron ejected from the nucleus of the atom Electrons do not exist by themselves within the nucleus- a neutron can split into two parts, becoming an electron and a proton. The electron is ejected and the proton stays behind in the nucleus. 14/6 C ------> 14/7 N + 0/-1 e (β) + (γ) Gamma Rays (γ): Gamma rays (γ) are a kind of radiation similar to light, microwaves, and x-rays except they are much higher in energy, so they can be very dangerous. When gamma rays are emitted, the identity of the emitting atom does not change. Gamma rays accompany alpha and beta decay and also fission. Half Life: Half life is the time required for a radioisotope to decay to half (½) of its original amount. Each radioisotope has its own half life. Half lives range from nanoseconds to billions of years. The shorter the ½ life, the more energetic (unstable)

Nuclear Fission and Fusion When the nuclei of certain isotopes are bombarded with neutrons, they undergo fission, the splitting of a nucleus into smaller fragments. In a chain reaction, some of the neutrons produced react with other fissionable atoms, producing more neutrons which react with still more fissionable atoms. Chain reactions can be controlled by: Neutron Moderation- slows the neutrons Neutron Absorption- decreases the number of slow moving neutrons Fusion occurs when nuclei combine to produce a nucleus of greater mass. In solar fusion, hydrogen nuclei (protons) fuse to make helium nuclei and two positrons. Fusion reactions, in which small nuclei combine, release much more energy than fission reactions, in which large nuclei split.

Chemical Vs. Nuclear Reaction Chemical Reaction: A new substance is created Electrons are transferred or shared between atoms (bonds) Small amounts of released energy Nuclear Reaction: A new element is created Occurs when nuclei combine or split (fusion, fission) Releases huge amounts of energy

Electron Shells Electron shells are the levels around the nucleus where electrons can be found The atomic number is the same as the total number of electrons The period number is the same as the number of electron shells For main group elements- the group number is the same as the number of valence electrons  Each shell has a maximum number of electrons Vocab: Valence shell- the outermost electron shell Valence electrons- electrons located in the valence shell Core electrons- all other electron besides the valence electrons Subshells: s, p, d, f s- max # of electrons = 2 p- max # of electrons = 6 d- max # of electrons = 10 f- max # of electrons = 14   Direction of increasing n+l value : 1s 2s   2p 3s   3p   3d 4s   4p   4d   4f 5s   5p   5d   5f... 6s   6p   6d ... ... ...

Ions Chemists have found that metal atoms transfer electrons to nonmetal atoms when they form compounds. Ions are formed when electrons are removed from or added to an atom. The rest of the atom stays the same (protons, neutrons, mass). The charge on an ion is noted with a superscript. Cations are positively charged (lose electrons)- usually formed from metal atoms Anions are negatively charged (gain electrons)- usually formed from nonmetal atoms When atoms gain or lose electrons, they form ions. Ions are atoms that carry a net positive or net negative charge. (When atoms lose electrons, they have a positive charge and are called cations. When atoms gain electrons, they have a negative charge and are called anions.) Ions have electron arrangements resembling those of the noble gas atoms. Atoms tend to lose or gain electrons to attain the electron arrangement of a noble gas. Atoms become ions to achieve the electron configuration of their nearest noble gas. In that way, atoms become more stable (lower energy state) Ionic Compounds: Ionic compounds form between metal atoms and nonmetal atoms. Electrons are transferred from one atom to another, forming a cation (+) and an anion (-). Cations (+) and anions (-) are attracted by electrostatic force (opposites attract). eg. Metal + Non-metal Every time a metal atom and a nonmetal atom bond, they form a compound with an overall zero charge. This is known as the rule of zero charge. A compound is always neutral. Because a compound needs to be neutral, you have to add the correct number of anions or cations to balance the compound to make it neutral. To name an ionic compound: Name the cation Add the anion change the ending to "-ide" Polyatomic Ions: Polyatomic ions contain more than one atom,. They are composed of a group of atoms with an overall positive or negative charge.. Most polyatomic ions are anions, with negative charges. The rule of zero charge can be used to predict the formulas of compounds that contain polyatomic ions.  Each polyatomic ion has its own name. Most polyatomics are indicated by the suffixes "-ite" and "-ate". Parenthesis indicate more than one polyatomic. When naming a compound, you insert the polyatomic ion name at either the beginning or ending of the chemical name. The cation is first and the anion is second. Transition Metals Transition metals have more than one oxidation state Charge is indicated by roman numerals (eg. Copper (II) Oxide) Except: Cd²⁺, Zn²⁺, and Ag⁺

Solubility and Conductivity dissolve- to disperse evenly into another substance (eg. solid dissolves into liquid) conductivity- a property that describes how well a substance can transmit electricity solubility- a property that describes how will a substance can dissolve into another substance Solubility and conductivity help to determine the types of bonds in a substance. Not all substances conduct electricity, and not all substances dissolve in water. Substances that do conduct electricity involve either solid metals, or metal-nonmetal compounds dissolved in water. By using solubility and conductivity tests, we can categorize matters into types of bonding. We can place all the substances into one of the four categories. They are also groups of bonding categories. Soluble and conducts Soluble and does not conduct Insoluble and conducts Insoluble and does not conduct

Types of Bonding A chemical bond is an attraction between atoms that holds them together in space. Atoms bond to achieve a more stable state. Bonds are formed between neighboring atoms by sharing or transferring electrons. Types of bonding: Ionic Electrons are transferred from one atom to another Made up of non-metal and metal atoms Soluble, Conductible when dissolved Usually brittle solids eg. salts (sodium/potassium/etc. chloride) Molecular Covalent Valence electrons are shared between pairs or groups of atoms, creating small stable units Made entirely of non-metal atoms Soluble (some), Non-Conductible Mostly liquids and gases eg. water (hydrogen oxide) Network Covalent Valence electrons are shared between the entire substance Made entirely of non-metal atoms Non-Soluble, Non-Conductible Extremely hard solids eg. sand (silicon dioxide) Metallic Valence electrons move freely about the substance Made entirely of metal atoms Non-soluble, Conductible Bendable, Malleable Solids eg. tin, gold, aluminum The HONC Rule: Hydrogen makes 1 bond Oxygen makes 2 bonds Nitrogen makes 3 bonds Carbon makes 4 bonds   Lewis Dot Structures Sh​​​​ows pairs of electrons, both bonded pairs and lone pairs lone pair- a pair of electrons not involved in bonding Atoms tend to follow the Octet Rule, except for Hydrogen, which follows the duet rule The number of dots correlates with the number of valence electrons Double and Triple bonds Single bond- shares 2 electrons Double bond- shares 4 electrons Triple bond- shares 6 electrons Electronegativity: Electronegativity- the tendency of an atom to attract shared electrons Metals- electron givers; lower electronegativity Non-Metals- electron takers; higher electronegativity Electronegativity increases from right to left across a period and up a group The difference between electronegativities determines how polar a molecule is

Chemical Formulas Empirical Formula- The formula of the compound with the smallest whole number ratio of atoms  Molecular Formula- Tells the number and type of atoms in a molecule Structural Formula- Shows the structure of a molecule and how the atoms are bonded or connected The smell of a compound is usually related to the structure of a molecule Molecules can have the same molecular formula, but different structural formulas- these are called isomers

Functional Groups Classification of Smells: Putrid- repulsive, gross Camphor- pungent, medicinal Sweet- flowery or fruity 🏳️‍🌈 Minty- green herbal smell Fishy- very distinctive, like seafood Organic Molecules contain a carbon backbone and one or more functional group Hydrocarbon- backbone for functional groups. They include: Alkanes- only single bonds Alkenes- contains double bonds Alkynes- contains triple bonds Functional group- a group of atoms bonded together in a specific way define the chemical and physical properties of a compound Types of Functional Groups: Alcohol Functional group: -OH Smell: Camphor/Medicinal Compound names end in: -ol Amine ​​​​​​​​​​​Functional group: -N Smell: Fishy Compound names end in: -amine Carboxylic Acid​​​​​​​ ​​​​​​​​​​​Functional group: -COOH Smell: Putrid Compound names end in: -ic acid Ketone ​​​​​​​​​​​Functional group: -COOC Smell: Sweet Compound names end in: -one​​​​​​​ Amino Acids ​​​​​​​Contain two functional groups- amine and carboxylic acids Building blocks of protein "R" group: represents a sidechain

Electron Domains The VSEPR (Valence Shell Electron Pair Repulsion) Theory states that molecules will attain whatever shape keeps the valence electrons of the central atom as far apart from one another as possible. Unshared pairs of electrons on the central atom determine the shape. Electron pairs REPEL each other.   Electron domain: The space occupied by valence electrons in a molecule, either a bonded pair(s) or a lone pair. Electron domains affect the overall shape of a molecule. Electron domain theory: The idea that every electron domain in a molecule is as far as possible from every other electron domain in that molecule. 4 bonding pairs, 0 lone pairs- tetrahedral shape (109.5 degrees) 3 bonding pairs, 1 lone pair- trigonal pyramidal (105 degrees) 2 bonding pairs, 2 lone pairs- bent 2 bonding pairs (double bonds), 0 lone pairs- linear (180 degrees) Lone pairs occupy more space than bonding electron pairs. Double bonds occupy more space than single bonds.

Polarity Static electricity When two different materials are rubbed together, some electrons can transfer This results in an imbalance of positive and negative charges Partial charges at different locations on a molecule cause a polar molecule Polar molecules are attracted to each other Polar Molecules: Molecules that are attracted to a charge because they have partial charges on them; one side is slightly negative and the other side is slightly positive Non-Polar Molecules: Molecules that are not attracted to a charge   Attractive molecules Partial Charges- molecules don’t share their electrons equally Polar molecules; have a dipole moment Non-polar molecules- share their electrons equally; have no charge 2 polar bonds may cancel out if they are going in opposite directions Intermolecular Forces- Attraction between molecules (not within) Forces that hold liquids and solids together When a substance melts or boils intermolecular forces are broken When a substance condenses or freezes, intermolecular forces are formed

Chemical Equations Chemical equation - a chemical “sentence” that describes change, using numbers, symbols, and chemical formulas A chemical equation can help you predict what you will observe when the reactants are combined Chemical equations describe chemical reactions Reactants ---------------------------------------------------------------------------------------------------------------> Products The starting materials of a chemical rxn                “yields"                                                             New substance that is formed Go through a chemical rxn                                        indicates the chemical rxn                           Listed on the right Listed on the left Example: H₂(g) + Cl₂(g) -----> 2HCl(aq) The subscripts show how many of each atom The coefficient shows how many of each molecule (s)- solid, (g)- gas, (aq)- aqueous, (l)- liquid-- show the state of matter Law of Conservation of Mass: Matter can be neither created nor destroyed in physical and chemical changes. Matter is conserved. Individual atoms are conserved in chemical reactions and physical changes, so the number of atoms remains the same from start to finish. Mass is also conserved, so the total mass of the products is equal to the mass of the reactants.    4 Steps for Balancing Equations: Get the skeleton equation Draw boxes around the chemical formulas Make an element inventory Update your inventory until balanced (# atoms of reactants = # atoms of products)

Physical vs. Chemical Change ​​​​​​Chemical reactions allow chemists to predict and track changes in matter. They indicate how many products are formed, what those products are, and the phase of each product.  Signs/Indications of a Chemical Reaction: Energy change (heat released or absorbed) Exothermic reaction- heat energy released Endothermic reaction- heat energy absorbed Production of gas Formation of precipitate Color change Change in odor or smell Physical Changes: Changes in the appearance or form of a substance. Phase changes (solid, liquid, or gas) When something changes in a physical way, the chemical formula does not change because you end up with the same substance you started with. Chemical changes produce new substances with new properties. New substances are formed during chemical changes, so we can expect the products to have properties significantly different from the reactants. Heat, light, smoke, smells, bubbles, and changes in color often accompany chemical changes. The terms chemical change and chemical reaction mean the same thing. DIssolving:  Chemical equations can also represent the process of dissolving.  eg: C₂H₆O (l) --> C₂H₆O (aq)        HCl (g) --> HCl (aq) Another way of writing the second equation: HCl (g) --> H⁺ (aq) + Cl⁻ (aq) Ionic compounds do not dissolve the same way as molecular solids. This can be shows with an equation that stresses the formation of ions in solution.           Molecular substances do not break apart when dissolved so there is only one way to write the equation Dissolving is generally considered a physical change, but it is common with chemical change as well

Types of Reactions Synthesis (or Combination) Reaction Two or more reactants (simple) to one product (complex) General Equation: A + X --> AX eg: 3H₂(g) + N₂(g) --> 2NH₃ Decomposition Reaction One reactant (complex) to two or more products (simple) General Equation: AX --> A + X eg: PbCO₃(s) --> PbO(s) + CO₂(g) Single-Replacement (Exchange) Reaction One element replaces another of the same type charge General Equation: A⁺ + B⁺X⁻ --> AX + B  eg: Cl₂(g) + KBr(aq) --> KCl(aq) + Br₂(g) Double Replacement (Exchange) Reaction Exchange cations and anions An acid and base react to form water General Equation: A⁺X⁻ + B⁺Y⁻ --> AY + BX eg: 2AgNO₃(aq) + MgCl₂(aq) --> 2AgCl(s) + Mg(NO₃)₂(aq) Combustion Reaction Substances combine with O₂-- must release heat or light Balance in order of C, H, then O General Equation: X + O₂--> CO₂ + H₂O eg: 2C₅H₁₀(g) +15O₂(g) → 10CO₂(g) + 10 H₂O(l)

Moles The Mole is the SI unit for amount of substance. It is used as a counting unit used to count a large number of atoms. 1 mol of atoms = Avogadro's number = 602 sextillion; 602,000,000,000,000,000,000,000 6.02x10²³ Molar Mass: The mass of 1 mol of a substance is called the molar mass Written in terms of grams per mole (g/mol) eg: Mercury (Hg) - 200.6 g/mol Avogadro's number: Is the number of particles in exactly one mole of a pure substance Can be written as: 6.022x10²³ particles/mol 6.022x10²³ atoms/mol 6.022x10²³ molecules/mol Mole Conversions:   Use the factor label method: Start with the end in mind List your given (put your given over 1) Use the appropriate conversion factor Check your units mass to moles: (# of grams/1)( 1 mol/Molar Mass (g)) moles to mass: (# of moles/1)(Molar Mass (g)/1 mole)

Solubility and Molarity Precipitation of ionic solids Sometimes, cations and anions combine and come out of the solution as a solid. This solid is called the precipitate, though precipitation are not limited to solids. Precipitate: a solid formed in a chemical reaction between two solutions Solubility: Solubility is the degree of which a compound dissolves in water. Ionic substances vary significantly in solubility, as some solids are more soluble than others. When a compound reaches the limits of its solubility, undissolved solid is visible. Important Terms: Solution: a mixture of two or more substances that is uniform throughout. Solute: substance dissolved in a solution (the stuff being dissolved) Solvent: substance in which the solute dissolves in a solution (the stuff doing the dissolving) Solution Concentration: The concentration of a solution is a measure of the amount of solute that is dissolved in a given amount of solvent. THis is called the molarity d a solution and is expressed in moles per liter, mol/L Molarity (M) = moles of solute/liters of solution If you want to make a solution more concentrated you need to add solute, and if you want to make a solution more dilute, or less concentrated, you should add water, or the solvent A solution is saturated when it can hold not more solute. The solution contains the maximum amount of solute for the given amount of solvent   If you know the molarity of a solution and the size of the sample, you can calculate the number of grams of substance dissolved in that solution: grams = moles x molar mass You can also use the "magic triangle":                    mols of solute (mol) Molarity (M)          Liters of solution (L) M= mols/Liters mols = Molarity x Liters Liters = mols/Molarity ***Molarity is always is units of mol/L, never mol/mL

Acids and Bases Acids and Bases Many acidic and basic solutions are colorless and odorless. There are molecular substances called indicators that change color when they come into contact with acids or bases. cabbage juice litmus  universal Acids: taste sour litmus paper ---> red react with strong bases to produce salt and H₂O Bases: taste bitter litmus paper ---> blue react with acids to produce salt and H₂O Substances with a pH 0-7 are acids, while substances with a pH 7-14 are called bases. Substances in the middle (7) are neutral.

The pH scale The pH scale is a logarithmic scale that describes the concentration of H⁺ ions in a solution. pH is related to [H⁺] by the formula: pH = -log[H⁺] In any solution, the product of the hydrogen ion, H⁺, concentration and hydroxide ion, OH⁻, concentration is a constant.    -Water dissociates into H⁺ and OH⁻ ions. The pH scale is a logarithmic scale that describes the concentration of hydrogen ions, H⁺ , in solution: pH = -log[H⁺]. The H⁺ concentration is related to the OH⁻ concentration: [H⁺][OH⁻] =10⁻¹⁴. As [H⁺] increases, [OH⁻] decreases, and vice-versa. The pH of water is 7. In water the H⁺ concentration is equal to the OH⁻ concentration. Thus, water is neutral.

Dilutions   Adding water to an acid or a base dilutes the solution, making it less acidic or less basic. Water added to acid: The [H+] concentration of a solution decreases, and the pH increases towards 7. Each time the H+ concentration is diluted tenfold, the pH number goes up 1 unit. Water added to base: The [OH –] concentration of a solution decreases, which increases [H+] concentration, and the pH decreases towards 7. ***An acid can never be turned into a base by diluting it with water, and a base can never be turned into an acid by diluting it with water. Diluting a solution reduces the number of moles of solute per unit volume, but the total number of moles of solute in solution does not change The total number of moles of solute remains unchanged upon dilution, so you can use this equation: M₁ x V₁ = M₂ x V₂ M₁ and V₁ are the molarity and volume of the initial solution, and M₂ and V₂ are the molarity and volume of the diluted solution.   Dilute solutions are less acidic and less basic than more concentrated solutions. • Dilution of an acid or a base results in a solution that is increasingly neutral. • It is not possible to go past a pH of 7 when diluting an acid or a base.

Neutralization Reactions and Titrations Neutralization: A neutralization reaction between an acid and a base in aqueous solution produces an ionic compound (generically referred to as a salt) and water. It can be described as a double exchange reaction in which the two compounds exchange cations. The pH of the product approaches 7. This is because the H⁺ and OH⁻ ions are being converted into water molecules. As their concentrations decrease, the pH will move closer to 7.    Titrations: A titration is a procedure in which a neutralization reaction is monitored with an indicator (such as phenolphthalein) allowing you to calculate the unknown concentration of an acid or base If the solution is acidic, a standard base solution is added until it is neutralized and vice versa. The end point is indicated by a color change and is when the number of H⁺ ions is equivalent to the number of OH⁻ ions. Calculations are based off of the following formula (Volume of acid)(Molarity of H⁺) = (volume of base)(Molarity of OH⁻)

Mol Ratios Balanced Equations show ratios usually in terms of moles.  2Al₂O₃ ---> 4Al + 3O₂ (For every 2 moles of Al₂O₃ that is used, 4 moles of Al and 3 moles of O₂ are produced)   Mole-Mole Problems: The quantity of one or more reactants is given in moles. The quantity of products in moles is requested. The mole ratio is used to calculate.   Excess and Limiting Reactants: If reactants are not combined in their mol ratio, one will run out and the other will be excess. The one that runs out is called the limiting reactant. The limiting reactant (or limiting reagent) in a chemical reaction is the substance that is totally consumed when the chemical reaction is complete. The amount of product formed is limited by this reagent as the reaction can not continue.  To determine the limiting reactant: Convert both the given masses (grams) to amounts in moles Then calculate the number of moles of one of the products The reactant yielding the smaller number of moles of product is the limiting reactant   Percent Yield % yield = actual yield/theoretical yield x 100% Theoretical Yield: is the maximum amount of the product that can be produced from a given amount of reactant Actual Yield: The measured amount of a product from a reaction

Exothermic and Endothermic Energy Transfer The chemical process you are focusing on is referred to as the system. The surroundings are anything outside of the system. A chemical process involves energy transfer due to temperature differences between two substances.  Endothermic (cold) transfer of heat to the system from the surroundings Exothermic (hot) transfer of heat from the system to the surroundings Thermodynamics Thermodynamics is the study of energy and its transformations  1st Law of Thermodynamics Energy is conserved- it doesn't just disappear, it is transferred 2nd Law of Thermodynamics Energy tends to disperse Entropy tends to increase Thermal Equilibrium- when the objects in contact are the same temperature You can not transfer energy from a cold object to a hot object because this would concentrate the energy

Specific Heat Temperature: If temperature is measures in Kelvin, then its value is directly proportional to the average kinetic energy of the molecules of a substance. Heat: Actual energy measured in joules or other energy units   Heat is a flow of energy due to a temperature difference Thermal Energy: Flow of energy called heat, transferred from a hot object to a colder object. Thermal energy refers to the amount of energy in a sample It is dependent on both mass and temperature of the sample Because heat is a process of energy transfer, it is incorrect to refer to the "heat energy of a sample" the correct term is "thermal energy of a sample" Heat Flow Heat flow is measured in two common units: the calorie and the joule The energy in food is expressed in kilocalories (kcal) 1 kcal = 1000 calories Units of Energy The calorie or joule is a unit of energy used to measure heat transfer and to express thermal energy It takes one calorie of heat to raise the temperature of one gram of water by one degree Celsius It takes 4.184 joules to raise the temperature of one gram of water by one degree Celsius. This is known as the specific heat of water Specific Heat Capacity There are two factors that the heat capacity of an object depend on: Mass (g) Chemical composition- what the material is made of Specific Heat Capacity Equation We use the following equation: q = m x C x ΔT q = heat (joules or calories) m = mass (grams) ΔT = change in temperature (°C)

Phase Change 6 Types of Phase Change Melting = Fusion (solid to liquid) Freezing = Solidification (liquid to a solid) Boiling = Vaporization (liquid to gas) Gas to Liquid = Condensation Solid to Gas = Sublimation Gas to Solid = Deposition   Specific Heat of Water Solid (ice): 2.1 J/g°C Liquid (water): 4.184 J/g°C Gas (steam): 2.0 J/g°C The heat (joules) added/removed during a phase change will break intermolecular forces if warming or allow intermolecular forces to form if cooling The freezing AND melting point of water is 0°C The boiling point of water is: 100°C Heat of Fusion Heat (joules) absorbed when 1 gram of solid turns into a liquid (melting) For H₂O, this is 334 J/g (6.02 kJ/mol) If you have 1 g of ice at 0°C it will absorb 334 joules of energy before it will melt into liquid water at 0°C The temperature remains constant Heat of Vaporization/Condensation Heat (joules) absorbed when 1 gram of liquid turns into a gas (vaporization) For H₂O, this is 2260 J/g (40.6 kJ/mol) If you have 1 g of water at 0°C it will give up 2260 joules of energy before it will solidify into ice at 0°C The temperature remains constant When we do problems that involve a phase change, we use Enthalpy q = mΔHfus (Melting or Freezing)  q = mΔHvap (Vaporization or Condensation)   Calorimetry Conversions 1.0 calorie = 4.184 J                        1 kilocalorie = 1000 calories                        1kJ = 1000 Joules

Gas Laws and Kinetic Molecular Theory Charles' Law The ratio of the volume and temperature of a gas is constant at a constant pressure. Since the ratio of the volume and temperature of gas is constant at a constant pressure, it remains the same at different times.  Equation:  V₁/T₁ = V₂/T₂   Boyle's Law The product of a pressure and the volume of a gas is constant at constant temperature. Since the product of the pressure and the volume of a gas is constant at constant temperature, it remains the same at different times. Equation:  P₁V₁ = P₂V₂   Gay-Lussac's Law The ratio of the pressure and temperature of a gas is constant at constant volume. Since the ration of the pressure and temperature is constant at constant volume, it remains the same at different times.  Equation: P₁/T₁ = P₂/T₂   Combined Gas Law The ratio of the product of pressure and volume to the temperature of a gas is constant.  Equation:  P₁V₁/T₁ = P₂V₂/T₂   Avogadro's Law  The ratio of the volume and mols of a gas is constant at constant temperature and pressure.  Equation:  V₁/n₁ = V₂/n₂   Ideal Gas Law The ideal gas law relates to the pressure, volume, temperature, and number of mols for a gas sample.  Equation: PV = nRT   Kinetic Molecular Theory Gas molecules are in constant motion. The gas molecules in a sample move at hundreds of miles per hour in straight line paths. Even though they are very tiny, the molecules inside the container hit each other and the walls of the container. There are so many collisions that they add up to a measurable pressure.  Gas pressure is caused by molecules hitting the walls of the container.  Pressure increases as the frequency of collisions increases. It also increases if the molecules hit the walls with greater force. Pressure and temperature are proportional if the volume of the gas does not change: P = kT. An increase in temperature increases both the collision frequency and the force with which the molecules hit the walls.