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Chemistry 113 Notes

by: Leah Nakaima
Leah Nakaima


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These notes cover what was on our first exam and part of what is going to be on our second exam.
General Chemistry 1
Class Notes
General Chemistry
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This 6 page Class Notes was uploaded by Leah Nakaima on Wednesday June 1, 2016. The Class Notes belongs to CHM 113 (LEC, LAB) at 1 MDSS-SGSLM-Langley AFB Advanced Education in General Dentistry 12 Months taught by Sreekaram in Summer 2016. Since its upload, it has received 3 views. For similar materials see General Chemistry 1 in Chemistry at 1 MDSS-SGSLM-Langley AFB Advanced Education in General Dentistry 12 Months.


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Date Created: 06/01/16
Chapter 1. Matter and Energy Matter Anything that occupies space and has weight, therefore it has volume and mass. States of Matter Solid;   Fixed shape and volume,   may be soft or hard, and rigid or flexible,   particles are close together and organized Liquid;   Vary in shape,   take the shape of the container,   particles close together but disorganized Gas;   No fixed shape or volume,   Particles farther apart and disorganized. Properties The characteristics that give each substance a unique identity.  Types of properties; Intensive properties.  They are properties that do not depend on the amount of materials present. They never  change (constant). They include boiling and melting points, densities, viscosity, etc. Extensive properties.  These are properties that depend on the amount of materials present. They change for  example masses, volumes, radii, heights etc. Changes Physical changes;  Particles before and after remain the same  No change in composition  For example liquids to solids or to gasses Chemical changes;  Particles before and after are different  Change in composition  For example if an electric current is applied on water, it breaks down to oxygen and  hydrogen gas. Energy It is the ability to do work. Energy is divided into two types i.e.  Potential energy; Energy due to position of an object. It is how high or low and object is.  Kinetic energy; Energy due to motion of an object. It is how fast or slow an object is.  Total energy is the sum of potential and kinetic energy. Energy changes. Energy is neither created nor destroyed. It is conserved and converted from one form to another. Note Unit Conversions are provided on the cover page of every chemistry paper. Common conversion factors Billion/ Giga (G) = 1000000000 Million/ Mega (1M) =1000000 Kilo (K) =1000 Milli (m) =0.001 Micro (µ) =0.000001 Nano (n) =0.000000001 Density Density = Mass/ volume At a specific temperature and pressure, the density of a substance is a characteristic physical  property and has a specific value. Significant figures They are recorded digits, both certain and uncertain. The greater the number of significant  figures, the greater the certainty of a number. While finding the number of significant figures,  Any zero In front of a non­zero digit is a place holder and not a significant figure.  Any zero between non­zero numbers is significant  If there is a decimal, any zero at the end of a number is significant  If there is no decimal, any zero at the end of a number is not significant. Example, 9823­4s.f, 0.04587­ 4s.f, 98045­5s.f, 65.850­5s.f, 2980­3s.f, 2.980­4s.f Rules in writing significant figures. Multiplication and Division. The answer takes the number of significant figures as the measurement with the lowest  number of significant figures. E.g. 9.2 x 6.8 x 0.3744 = 23­2s.f. Addition and Subtraction. The answer takes on the number of decimal places as that of the measurement with the  fewest number of decimal places. E.g.2.5 + 0.32 = 2.8 ­ 1d.p Exact numbers, don’t have any un­certainty and do not limit the number of significant  figures in a calculation. Precision, Refers to how close the measurements are to each other in a series. E.g. 80, 81, 82, 78, 79­ are precise. 1, 10, 30, 19, 80­ Not precise Accuracy, Refers to how close the measurement is to the actual value E.g. if the boiling point is 100°c, 101°c is accurate, and 42°c is not accurate. Chapter 2 ­ Atoms, Ions and Periodic Table. Atomic Theory of Matter  The theory that atoms are the fundamental building blocks of matter reemerged in the  early 19  century championed by John Dalton.  Each element are made up of small particles called Atoms.  Atoms cannot be broken down into smaller particles either by physical or chemical  means. Law of conservation of Mass.  The total mass of substance present at the end of a chemical process is the same as the  mass of substances present before the process took place.  Thompson brought forward the “plum pudding” model­which was proved incorrect.  The bread was the positive charge while as the plums were the negative electrons.  Ernest Rutherford shot a particle at a thin sheet of gold foil and observed the particles  scatter within the particle. It postulated a very small dense nucleus with electrons around  the outside of the atom. Subatomic particles Protons­ Rutherford 1919 – positive (+) charge­ 1.0073 Neutrons­James Chadwick 1932 ­ No charge­ 1.0087 4 Electrons­ negative (­) charge­ 5.486x10 ­smallest.  The first two makeup the nucleus. Symbols of elements. Elements are symbolized by one or two letters. The mass number that appears at the top of a symbol is the sum of protons and neutrons The atomic number that appears at the bottom of a symbol is the number of protons or  electrons of the element. Proton number (p) is the same as number of electrons (e). Periodic Table.  Elements are arranged in order of their atomic masses developed by Dmitri Mendeleev.  Rows are periods and columns are groups.  Elements in the same group have similar chemical properties  Metals are on the left side and nonmetals are on the right side of the periodic table  besides only hydrogen. Metalloids border the stair­step line. Ions Atoms lose or gain electrons to form ions. Anions  Are negatively charged ions,   Formed by nonmetals on the right side of the periodic table,   The more electrons gained, the more negatively charged an ion becomes. Cations.  Are positively charged ions,   Formed by metals on the left side and middle of the periodic table.   The more electrons lost, the more positively charged. Monoatomic is one ion, and polyatomic is more than one ion. Seven molecules occur naturally as diatomic in sequence of hydrogen, nitrogen, fluorine,  oxygen, iodine, chlorine and bromine. Isotopes. They are formed when an atom gains or losses a neutron. For example carbon has three  isotopes i.e. carbon­12, carbon­11 and carbon­13.  Chemical formulas. Consist of element symbols with numerical subscripts. Subscripts on the right show the  number of atoms of that element. Types of compounds Ionic compounds.  Involves transfer of electrons from metals to nonmetals  Crystal formed which is tightly packed.  One of the strongest bonds  High melting and boiling points for example 1000°c  Solid at room temperature e.g. NaCl where Na is a metal and Cl is a nonmetal. Writing formulas for ionic compounds,  Write cation with a positive charge and the anion with a negative charge  Crisscross the numbers and if the subscript is not in the lowest whole number ratio,  simplify the numbers. Naming Ionic compounds.  First look at cation, write the name of the cation (Na = sodium) and If the cation can have more than one possible charge­transition metals, write the charge as a Roman numeral in  parentheses e.g. copper (II)  Then look at anion, if it is an element, change its ending to –ide (Cl = chloride), and if it  is a polyatomic ion, simply write the name of the polyatomic ion. (NO3 = nitrate)  DO NOT use prefixes di, tri, tetra even if there are multiple anions/cations Covalent (molecular) compounds  Elements share electrons and covalent bonds are formed between nonmetals.  Atoms far apart don’t interact, and for those closer, attractions between electrons of  different nuclei increases.  Not as strong as the ionic bond.  Liquids and gases at room temperature e.g. water  Low melting and boiling point (50°c­100°c) Naming of covalent compounds.  For the first element, write the prefix and then the name. Mono is not used  For the second element, write the prefix and change the ending element to ide Note. Compounds containing ammonia are not covalent for example NH Cl and N4 NO  are  4 3 ionic because they can breakdown into ions hence they are ionic. Molecular masses from chemical formulas. Molecular mass is the sum of atomic masses. For example for H O2   Molecular mass =  (2 x atomic mass of H) + (1 x atomic mass of O) Note. Masses are read from the periodic table.


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