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Chemistry Week 1 Full set of notes lectrues 1-3

by: Tiffani Marie

Chemistry Week 1 Full set of notes lectrues 1-3 CHMY 141

Marketplace > Montana State University > Chemistry > CHMY 141 > Chemistry Week 1 Full set of notes lectrues 1 3
Tiffani Marie

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Chemistry lectures on atoms
College Chemistry I
Bahn, Christian S
Class Notes
atoms, anion, ion, cation, Protons, neutron, Chemistry, Mixtures, elements, atomicmass, periodic table
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This 7 page Class Notes was uploaded by Tiffani Marie on Friday September 2, 2016. The Class Notes belongs to CHMY 141 at Montana State University taught by Bahn, Christian S in Fall 2016. Since its upload, it has received 202 views. For similar materials see College Chemistry I in Chemistry at Montana State University.


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Date Created: 09/02/16
Week 1 Day One Lecture One 8/30/2016 Chemistry is considered to be the central science behind the question. Chemistry involves the  study of matter and its transformations. Matter  Matter is “stuff” matter is anything having mass or that takes up space. Most matter is  easily identifiable in categories such as solid, liquid, or some not as easily identified  gases such as air, CO, CO2, and H2S. During any type of physical or chemical  transformation the total mass of the matter involved cannot change.   Matter can be divided in to mixtures and pure substances.  o Pure substances then can be divided in to elements and compounds.   Elements consist of one type of atom and are given an elemental symbol  such as gold.   Compounds however, are classified as two or more types of atoms that are bonded and have a composition given by a formula like water – H2O.  Molecules are the simplest form of a compound, unlike an element where  the simplest from are atoms (multiple bonded molecules). o  Mixtures can be divided in to two categories, either Homogeneous or  Heterogeneous.   Homogeneous mixtures have the same composition throughout the  sample.   Heterogeneous mixtures have different composition throughout the  sample, if you can visibly tell a difference it is heterogeneous.   In order for atoms to bond there must be a chemical reaction otherwise you just have two  pure substances. Identifying Types of Substances and Mixtures  Are pure substances separable into simpler substances?   If yes, it is a compound.  An example of a compound would be substance would  be water.  If no, it is an element. An example of an element would be Helium.  Is the mixture uniform throughout?   If yes, the mixture in homogeneous. Things like tea and sugar mix thoroughly and distribute evenly making it one homogenous mixture.    If not, the mixture is heterogeneous. something like wet sand has two different  types of particles that separate into distinct regions (i.e. The sand and the water).  The Scientific Method The Scientific Method is a never ending cycle of constant testing, observing, and refining  scientific laws. 1. Begins with an observation or a question 2. A hypothesis is created based on observations 3. Hypothesis is tested repeatedly to confirm theory or present a new hypothesis 4. A series of confirmed hypotheses lead to the creation of a scientific law Day Two 8/31/19 Measurements  Scientific observations can be qualitative or quantitative  o Qualitative is useful and can describe things like color o Quantitative uses units and has a precise amount or measurement  In order to give meaning to quantitative measurements we assign units of measurements  The most common system of measurement used in the sciences is the SI system or the  metric system.  o Length = Meter o Mass = Kilogram o Time = Second o Temperature = Kelvin (uses absolute zero and never negative making equations  easier) o Amount of substance = Mole o Electric current = Ampere o Luminous intensity = Candela Lecture Two Si Modifiers  These are prefixes that are placed in front of the base unit in order to express values  larger or smaller than the base unit  Know how to use and convert the most common o Nano (n)  o Micro  o Milli (m) o Centi (c) o Kilo (k) o Mega (M) Dalton’s Theories  Considered the father of modern chemistry  Started looking at ideas proposed by Democritus o Smallest possible particle ­ atom  Three very important laws of matter expanded out of Dalton’s research Conservation of Mass  This law states that matter is neither created nor destroyed o We can change the form of something but the atoms that were present at the start  are present at the end of the reaction o Only applies to chemical and physical reactions (phase change ie liquid to solid) o Total mass in a reaction must be conserved  Initially, this was a difficult law to prove because evidence seemed to exist to prove it  wrong  o Mass of reactants = Mass of product Law of Definite Proportions  All samples of a given compound always have the same mass proportion of their  elements o Ie – can’t have water with extra oxygen o This also means that chemical formula for the same compound will always be the  same o Pure water always 88.8% oxygen, 11.2% hydrogen (By mass) Law of Multiple Proportions  When two elements can combine in multiple ways, if 1g of element A is combined with  element B, the ratio of the masses of element B reacted will be a small whole number.  Dalton’s Atomic Theory  Each element is composed of tiny, indestructible particles called atoms (No longer true,  atoms are destructible)  All atoms of a given element have the same mass and same properties (same properties is 99.99% true, however mass is not true due to isotopes)  Atoms combine in simple whole number ratios to form compounds (Law of Multiple  proportions)   Atoms of one element cannot change into another element  The Electron  J.J. Thomson was working with cathode rays o Really just a beam of electrons (but he didn’t know that)  Beam could be deflected with an electric or magnetic field o Had a negative charge o Measured charge to mass ratio and realized particles had a mass ~ 1/ 2000  of a  hydrogen atom  Particle was called the electron st o 1  subatomic particle found Milliken’s Oil Drop Experiment  Robert Milliken determined the charge on an electron o ­1.6 x 10^­19 C (Colom – massive charge)  o By combining this value with Thomson’s charge­to­mass ratio the mass of the  electron was determined o Mass = 9.1 x 10^­28g Early Atomic Structure  Once the electron was discovered, models began to be proposed to fit the new data  Thomson proposed the plum pudding model (looks similar to a rasin cookie) Rutherford’s Experiments  The plum pudding model was not particularly robust model with lots of issues  1900’s saw an increase in radioactivity research  Experiments involved shooting a beam of alpha­ particles at gold foil and seeing what  happens Modern Nuclear Theory   Rutherford’s experiment led to a new theory about the atom o Most of an atom’s mass in concentrated in a nucleus o Most of the volume of an atoms is empty space o There must be a negatively charged particle in the center of the mass to counter  the negative electrons ­ Protons Day 3 09/02/16 Lecture 3 Modern Nuclear Theory  Rutherford’s led to a new theory about the atom o Most of an atom’s mass is concentrated in a nucleus o Most of the volume of an atom is empty space o There must be a positively charged particles in the center of the mass (nucleus) to  counter the negatively charged electrons – Protons this is because all atoms are  electrically neutral Neutrons  Even with the discovery of the proton, scientists knew the atom was not complete  o Relative masses could not be calculated correctly o Hypothesized missing mass was in nucleus  James Chadwick discovered the neutron in 1932 o Difficult to isolate due to lack of charge o Mass of neutron and proton are nearly identical  Both are similar to the mass of an H atom Subatomic Particles  Masses are often expressed in atomic mass units (amu) o 1 amu = 1/12 the mass of a Carbon­ 12 atom  Charges are often expressed in relative units o Charge of electron is ­1 o Charge of proton is +1 Elements  All elements are composed of the same three subatomic particles o All subatomic particles are the same as any other from different elements  Ie – a carbon proton is the same as an oxygen proton  So what makes one element different from another?  The number of protons in the nucleus is what defines the element o This number is called the atomic number and is noted as ‘Z’  The periodic table is arranged by atomic number Periodic Table  All elements are defined by a 1 or 2 letter symbol (3 for unnamed elements) o The atomic number listed on top tells us the number of protons and vice­versa  Some symbols are intuitive (He for Helium)  Some are not as easy to figure out (Na for Sodium) o However, Natrium is Latin for sodium Isotopes  Not all atoms of the same element have the same mass o All atoms of the same element have the same number of protons o Some atoms of the same element have a different number of neutrons  o We call these isotopes of each other  Isotopes may be man­made or naturally occurring (like C­12)  Most elements have isotopes but few do not The language of Atoms  The atomic number, meaning number of protons, is ‘Z’  We also have the mass number noted as ‘A’ o A can be found by taking Z, the number of protons, and adding the number of  neutrons o A = Z + #n  We can write the full atomic symbol for an isotope, giving us a lot of information about  the atom  We also refer to different isotopes by the symbol and mass number o An example would be C­12, which is read as “Carbon twelve” Ions  An ion is an atom that has obtained a charge o Either + or ­  The gaining of a charge cannot happen through loss or gain of protons o Nuclear particles are held too tightly in stable atoms  Charges are formed by the loss or gain of electrons o A loss of an electron(s) will lead to a positive ion called a cation o A gain of an electron(s) will lead to a negative ion called an anion Atomic Mass  How do we report the mass of an element when all the isotopes of an element have  different masses? o We use a “weighted average” of the naturally occurring isotopes  Based on the percentage of each isotope  Chlorine is a good example o Chlorine has two natural isotopes; Cl­35 (75.77%) and Cl­37 (24.23%)  Atomic mass = 0.7577(34.97amu) + 0.2423(36.97 amu) = 35.45 amu


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