CHEM 1030 notes week 3
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This 6 page Class Notes was uploaded by Alyssa Anderson on Sunday January 31, 2016. The Class Notes belongs to CHEM 1030 at a university taught by Dr. Streit in Spring 2016. Since its upload, it has received 56 views.
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Date Created: 01/31/16
Alyssa Anderson Chemistry Notes Week 3 An atom is the smallest quantity of matter that still retains the properties of matter An element is a substance that cannot be broken down into two or more similar substances by any means (such as gold, oxygen, helium) Atoms can also be divided smaller and smaller and eventually only a single atom remains. Dividing it further would make pieces that are no longer atoms Dalton- said atoms (of which matter consists of) are tiny, invisible particles Once a single atom has been obtained, dividing it smaller produces subatomic particles The nature, number, and arrangements of subatomic particles determine the propertied of atoms NOTE: like charges repel each other, opposite charges attract JJ Thompson (1856-1940)- noted easy were repelled by a plate with a negative charge and attracted to a plate bearing a positive charge. He concluded the rays were negatively charged. His contributions include: 1. Proposed rays were actually a stream of negatively charged particles 2. Negatively charged particles equaled electrons 3. By varying the electric ﬁeld and measuring the degree of deﬂection of cathode rays, Thompson determined the charge-mass ratio R.A. Milikan (1868-1953)- determined the charge on an electron by examining the motion of tiny oil drops The charge was found to be -1.6022 x 10^-19 C The mass of an electron equals the charge divided by the charge multiplied by the mass which means (-1.6022 x 10^-19)/(-1.76 x 10^8 C x grams) which means it equals 9.10 x 10^-28 grams Wilhelm Rotgen (1845-1923)- discovered x-rays which are not deﬂected by magnetic or electric ﬁelds so that they could not consist of charged particles Antoine Becquerel (1852-1908)- discovered radioactivity Alpha rays- consist of positively charged particles called alpha particles (α) Beta rays- electrons that are deﬂected and made of beta particles (β) Ernest Rutherford- used α particles to prove the structure of atoms 1. The majority of particles penetrated the gold undeﬂected 2. Sometimes, a gold particle would be deﬂected at a large angle or even backwards 3. Through this, Rutherford concluded the nuclear model which states a positively charged center is concentrated in the middle of a cell at the nucleus and that the nucleus accounts for most of the cell’s mass and is extremely dense at the core within the atom Protons- positively charged, in the nucleus Neutrons- no charge, in nucleus, slightly larger than protons Electrons- negatively charged particles that orbit around a nucleus BE SURE TO LOOK AT TABLE 2.1 IN THE BOOK All atoms can be identiﬁed by the number of protons/neutrons they have Atomic Number- number of protons in the nucleus 1. Since atoms must stay neutral, the number of protons equals electrons 2. Protons determine the identity of the element Mass Number- number of protons added to neutrons Most atoms have at least two isotopes, which mean they have the sam amount of protons and electrons, but different number of neutrons, which effects the mass. They usually exhibit the same chemical properties, such as some have the same type of compound with similar reactivities. On occasion, an isotope will be radioactive. Nuclear Stability- can be related to density (note: the total volume is hardly accounted for by the nucleus but the mass is mainly the nucleus alone) 1. The higher the density, the stronger the forces are in the atom. 2. Stability = Coulomb repulsion - short range attraction 3. example: the atomic number of Uranium equals 92 (protons and electrons) but 143 neutrons 4. Heavy atoms need much more neutrons to remain stable 5. The principle factor for nuclear stability is proton to proton ratio (n/p) 6. There are more stable nuclei with 2, 8, 20, 50, 82, or 126 protons and neutrons 7. There are more stable nuclei with even numbers 8. All elements with atomic numbers greater than 83 are radioactive LOOK AT TABLE 2.2 IN BOOK Atomic Mass- mass of atom in amu (1 amu = half the mass of a carbon-12 atom) Average atomic mass- on the periodic table, it represents the average mass of the naturally occurring mixture of isotopes The Periodic Table- a chart in which elements having chemical and physical properties are grouped together 1. Numbered by increasing atomic number (protons and electrons) because it regulates all properties of that element (ﬁngerprint of element) 2. There are two important numbers- the average atomic mass and the atomic number The periodic table is separated by periods and groups 1. Periods- horizontal rows, in order of increasing atomic number A. Metals- good conductors of heat and electricity B. Nonmetals- poor conductors of heat and electricity C. Metalloids- intermediate properties 2. Groups- vertical columns A. Alkali Metals (1A)- Li, Na, K, Rb, Cs, Fr B. Alkaline Earth Metals (2A)- Be, Mg, Ca, Sr, Ba, Ra C. Chalcogens (6A)- O, S, Se, Te, Po D. Halogens (7A)- F, Cl, Br, I, At E. Noble Gases (8A)- He, Ne, Ar, Kr, Xe, Rn F.Transition Metals (1B and 3B-8B) Mole- the amount of a substance that contains as many elementary entities as there are atoms in exactly 12 grams of carbon-12 1. Experimentally determined number- Avagandro’s Number (N ) A 2. NAaka 1 mole = 6.0221415 x 10^23 (usually simpliﬁed to 6.022 x 10^23) 3. example: The human body has a total of 30 moles of calcium. Determine the number of atoms of calcium and the number of moles of Calcium in a sample containing 1.00 x 10^26 Ca atoms. work: (30 moles Ca) x (6.022 x10^23/ 1 mole Ca) = 1.807 x 10^25 atoms of Ca Molar Mass- mass in grams of 1 mol of substance 1. By deﬁnition, the mass of one mole of carbon-12 is exactly 12 grams 2. The mass of 1 carbon-12 atom equals exactly 12 amu 3. The mass of an atom equals the mass of the mole (just in different units) 4. example: determine the number of moles of carbon in 25 grams of carbon and the number of moles of He in 10.50 grams of Helium work: (25 grams of C) x (1 mole of C/ 12.01 grams of C) = 2.082 moles of C (10.50 grams He) x (1 mole He/ 4.003 grams C) = 2.633 mol He 5. example: (0.515 g C) x (1 mol C/ 12.01 g C) x (6.022 x 10^23/ 1 mole C) = 2.58 x 10^22 C atoms Unit of energy is the Joule (J) The Joule was created by English physicist James Joule It is the amount of energy possessed by a 2 kg mass moving at speed of 1 m/s E K 1/2 x m x u^2 = 1/2 x 2 kg x (1 m/s)^2 = 1 kg x m^2/s^2 Joules can also be denied as the amount of energy exerted when a force of 1 Newton is applied over 1 meter. 1 J = 1 N x m The nature of light- visible light is only a small component of the continuum of radiant energy known as the electromagnetic spectrum Properties of waves (all forms of electromagnetic radiation travels in waves) 1. Waves are characterized by wavelength (λ) which is the distance between identical points on successive waves 2. frequency (v, nu) is the number of waves that pass through a particular point in 1 second 3. NOTE: short wavelength = high frequency; long wavelength = low frequency 4. v is inversely proportional to λ 5. amplitude is the vertical distance of the middle of the wave 6. The speed of light (c) through a vacuum is constant. A. c = 2.99792458 x 10^8 m/s B. y is proportional to x, so y = x 7. The speed of light, frequency, and wavelength are all related by the equation c = λ x v (λ is expressed in meters, and v is expressed in s^-1)
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