CHM110 - Chemical Principles 1

Nanomaterials, Nanotechnology, Nanoparticles Macroscopic - everything we can see (millions of molecules together) Microscopic - single molecules Nanoparticle - contain 1000 atoms or molecules and can be seen under TEM Property changes when becomes nanoparticle.

Unit One: Lecture Notes Atoms # (protons) #e (electrons ) # atomic number ## mass number For equal # the # identifies the isotope Atomic Mass Mass Number Mass of atom (amu) Mass of atoms (g) amu/g Mole particles Shorthand Notation for the Atom Components of the Atom Percent Abundance (where is the % abundance of ) Molecules Formed when e- of one atom simultaneously is attracted to its own nucleus and another surrounding nucleus. Example When ➀ and ➁ move closet, attracted

Units Significant Figures - Refers to digits that were measured when rounding calculated numbers we pay attention to significant figures, so that we do not overstate the accuracy of our answers. Example - Better measurement, more accurate - Less accurate Accuracy - Proximity of a measurement to the true value of quantity. Precision refers to proximity of several measurements to each others. What Exists? Atoms Atomic Number - Number of protons in the atom Mass Number - Protons a

Micro - Only be seen with a microscope Macro - Can be seen with bare eyes Volume and - most commonly used Uncertainty in Measurement Use burrette for accurate (to the decimal) measurements Use graduated cylinder for precise (rounded) measurements Accuracy vs Precision What Exists? Atoms number of protons = number of electrons = atomic number number of protons + number of neutrons = mass number mass of 1 atom isotopes Atoms of the same element with different number of neutrons =

Boyle’s Law Pressure of a fixed mass of gas is inversely proportional to its volume at constant temperature. Initial Conditions Final Conditions at constant and Why? Constant value because moles and temperature are constant Proportionate to volume If volume increases, then as is constant Charle's Law The volume of fixed mass of gas is directly proportion to it’s Kelvin temperature at constant pressure Initial Conditions Final Conditions at constant and

Ex: Chlorine has two isotopes and . is 75%. What is the atomic mass of ? Ex: Uranium has atomic mass of 238.03 and it consists of 2 isotopes with mass of with mass of What is the % abundance of in naturally occurring uranium? What Exists? Matter Pure - Constant composition Mixtures - Variable composition Homogeneous (single phase) Heterogeneous (more than once phase)

Laws of Thermodynamics 1. Where: hear absorbed by the system work done on the system 2. for all spontaneous processes 3. entropy of a perfect crystal at Enthalpy change for a process Heat absorbed by the system when the process is carried out at constant pressure final stateinitial state If Exothermic if Endothermic Calculating/Measuring Carry out the reaction in an isolated system Measure T due to reaction Calculate H from & What is ?

Nucleus Binding energy released when nucleus is formed Nucleic energy A large amount of energy is produced by binding energy 1. Example Chlorine has 2 isotopes: 3s Cl and 37 Cl. Cl is 75%. What is the atomic mass of Cl. 35x75 = 262.5 37x25=935 => 925+262.5 =35.5 100 2. Example Uranium has atomic mass of: 238.03 It consists of 2 Isotopes 235: U with mass of 235.044, and 238U with mass of 238.05^ X: % of 238 U 238.051* X + (100-x) 235.044 = 238.03 * 100 238.051X + 2350YYX

Separation Techniques of mixtures Qualitative analysis Separation and identification of the components of the mixture, what is present. Quantitative analysis Determination of the quantity of each component of the mixture. how much is present. Crystallization - Separation of solids from a liquid solution. A crystal is an ordered arrangement of molecules. First formation of a crystal needs another material for the crystallization to happen (even dirt). Evaporation of water for example to form a N

Qualitative - Description with words Quantitative - Quantity shown with numbers Separation Techniques Crystallization Distillation Chromatography more affinity to the solvent travel more/faster/eluted first highest on the paper Solvent extraction Calculating Moles Mole is a number # of mols = mass in g Ex: How many atoms are in of copper? Ex: How many grams of will precipitate out if reacted? Combustion Reaction Combustion - Substance reacts with to form and Ex: Upon comb

Compressibility (P, V, T) Boyle's Law: (constant T and n) Charles Law: (constant T and n) Gay-Lussac: Ideal Gas Law more collision between molecules at different pressure relationship between and change, gas constant Kinetic Molecular Theory Molecules are far apart; most of the volume is empty. Molecules are in constant random motion they travel in straight line paths. Molecules exert no force on each other except during collisions. Average kinetic e

Find the Empirical Formula of Vitamin C Example A 6.49 mg sample of ascorbic acid which is known to contain only and was burned in a combustion apparatus. The increased weights of the absorption tubes showed that 9.74 mg of and 2.64 mg of were formed. What is the empirical formula of the acid? 6.49 mg 9.74 mg 2.64 mg Moles of Mass of compound = number of moles atomic mass # moles of Find the % Composition of a Mixture Example A mixture of and with a mass of 4.000

Moles Units mole of oxygen atoms = atoms = mass of grams mole of carbon atoms = atoms = mass of grams mass of carbon atom = Calculation of Moles How many moles in 36g (C) Carbon? # How many moles of NaOH are there in 8g? = Calculate the # of atoms in mol of Copper. How many atoms are there in g of Carbon? A silver ring contains silver atoms. How many moles of silver are in a ring? # How many grams of Ag Cl will precipitate out if mole is reac

Solubility Rules What Happens? Reactants and products? number of reactants number of products Elements and compound Neutralization/acid-base reactions Acid - Produces from Base - Produces from Solubility Guidelines soluble Halides insoluble Most salts of and Important Redox Reaction Oxidation - Loss of electrons Reduction - Gain of electrons Oxidizing Agent - Causes an element/compound to lose electrons and gets reduces Reducing Agent - Causes an element/compound to

Principles of Chemical Equilibria Definition Rate - Change in [ ] / time at equilibrium K - equilibrium constant k - rate constant - # moles of gaseous product - # moles gaseous reactant Factors Affecting 1. temperature Exothermic : , (shift left) , (shift right) b. Endothermic : , (shift right) , (shift left) Van Hoff Equation assume derived from it, NOT the equation Quantitative Exothermic (negative #) as

Combustion - A reaction in which a substance reacts with oxygen to form oxygen containing compound and heat. Example - Combustion of natural gas (methane) produces carbon dioxide and water Empirical Formula = Molecular Formula Ratio of atoms in a molecule - Actual number of atoms molar mass/empirical formula molar mass Example 1 - Aspirin Given - % % % Find - Empirical formula Moles of each element - Mass of Mass of Mass o

DEFINE a) ACIDS AND BASES 1. arrhenius (has to be in water) Acid - proton donor, will donate , yields Base - yield when dissolved in 2. bronsted-lowry (does not have to be in water) Acid - proton donor, donate Base - proton acceptor 3. lewis Acid - accepts pair of Boron will accept pair of Base - donates pair of can donate that pair to make a bond 4. solvo-system Acid - Generates the cation form of solvent in solution Base - Generates the anionic form of solvent i

Balancing Reactions Balancing Just add Example removed since remained throughout the reaction Example

Test for Present Chloride Using Silver nitrate, if there is white precipitation, there is Chloride. precipitation reaction mass % of each component of original mixture mass mass of mass mass of Solve for Mass of divalent = Calcium (by the valency) Metal Chloride = % of Chemical Reactions 1. Classify by number of reactants and products A. Number

Concentrations of Solutions % w/w - g solute/100 g solution % w/v - g solute/ 100 ml solution % v/v - ml solute/100 ml solution Example 20 ml of a 5% solution of can completely neutralize 20 ml of a 5% solution of . True or False? False → Since we don't know if the number of mole is equal to each other or we don't know how many moles are involved. molar solution Concentration expressed in mol Molarity (M) = moles of solute / L solution Example What is the molarity of a 40 moles of dissolve