General Chemistry 1315 Notes Week 2
General Chemistry 1315 Notes Week 2 Chem 1315-003
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Popular in Chemistry
This 11 page Class Notes was uploaded by Ryan Henry on Friday January 29, 2016. The Class Notes belongs to Chem 1315-003 at University of Oklahoma taught by Dr. Awasabisah in Spring 2016. Since its upload, it has received 35 views. For similar materials see General chemistry in Chemistry at University of Oklahoma.
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Date Created: 01/29/16
General Chemistry Notes Week 2 Introductory Vocabulary ● Matter Anything that has mass and occupies space ● Massa quantitative measure of the amount of matter in an object ● Weight is a measure of the force of gravitational attraction between an object and a significantly large body (mass x acceleration due to gravity) ● Matter ● Elements a substance that cannot be separated into simpler substances by chemical means. ● Atom The smallest representative part of an element Atoms are the building blocks of matter Atoms combine to form molecules Molecules are a chemical combination of 2 or more atoms (ethyl butanoate pineapple) (methyl butanoate apple) With the addition of one carbon and two hydrogen atoms, these two substances will have very different odors/flavors. The atomic theory of matter ● Law of conservation of mass In a chemical reaction, matter is neither created nor destroyed. Atoms of an element are not changed into atoms of a different element by chemical reactions. The combined masses of the reactant(s) in a chemical reaction must equal the combined masses of the products. ● ● Law of Definite Proportions/Law of Definite Composition: All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements.If two elements such as the carbon and oxygen shown (left) combine to form two different compounds, the mass of oxygen that can combine with 1 g carbon can be expressed as a ratio of small whole numbers. EX: Decomposition of 18.0 g of water results in 16.0 g oxygen and 2.0 g hydrogen ● John Dalton’s atomic theory (1808): Each element is composed of extremely small, indestructible particles called atoms. Each element is composed of only one type of atom. The current model of the atom was developed through experiments that examined the behavior of charged particles. ● Particles with the same charge repel one another. ● Particles with opposite charges are attracted to one another. J. J. Thomson: Rays (electrons moving from the negative electrode to the positive) were the same regardless of cathode material – he discovered the electron.Constructed a cathoderay tube with a fluorescent screen/grid at one end in order to quantitatively measure effects of electric and magnetic fields ● Thomson was able to calculate a chargetomass ratio of the electron, 1.76 x 108 coulombs per gram. Millikan: Devised experiments to determine an experimental value for the charge of an electron and then discovered electron mass = 1.60 x 1019 C = 9.10 x 1028 g 1.76 x 108 C/g Ernest Rutherford performed an experiment (Gold Foil Experiment) to examine the angles at which alpha particles (positively charged particles) were deflected as they passed through a piece of thin gold foil. ● Most surprisingly, some particles scattered back in the direction from which they came (major deflection noted). ● Rutherford postulated that most of the mass of an atom resides in a small, dense, positively charged region he termed the nucleus. ● Rutherford further postulated that most of the total volume of an atom is empty space. Protons, discovered in 1919 by Rutherford, are positively charged subatomic particles. Neutrons, discovered in 1932 by James Chadwick, are neutral subatomic particles Location of subatomic particles within an atom: ● Protons and neutrons reside in the nucleus of the atom. ● The area surrounding the nucleus is a diffuse region of negative charge where the electrons reside (more detailed in Unit 4). Electrostatic repulsions between the positively charged protons should repel each other, but the nucleus does not fly apart because of the strong nuclear force, an extremely powerful but very short range attractive force between protons and neutrons. Characteristics of an atom: ● An atom must have the correct ratio of protons to neutrons for stability. ● A stable atom is one that does not have changes in the protons or neutrons of the nucleus. ● Too many or too few neutrons causes an atom to be radioactive, unstable. ● A change in the makeup of a nucleus, is a nuclear reaction – not a chemical reaction. 14C is an unstable atom. The ratio of protons to neutrons isn’t ‘right.’ ● Isotopes have a different mass number and are therefore unstable What about the number of neutrons in an atom? ● Atoms within an element (same atomic number) may have varying numbers of neutrons in the nucleus. ● While neutrons do not possess a charge, they have mass, therefore contribute to the overall mass of an atom. Ions: Ions are electrically charged atoms or groups of atoms called polyatomic ions. ● Ions may be positively charged or negatively charged depending upon whether the atom has lost or gained electrons. Ions are represented by using the chemical symbol for the atom with a notation of the charge as a superscript to the upper right of the symbol. Cations are PAWSitively charged Anions are negatively charged Atoms have very small masses – the atomic mass unit (u) is used instead of describing mass in absolute grams, i.e. 1.67x1027 kg. Early experiments could identify the relative masses of elements that composed substances. ● Water appeared to contain 8 times as much oxygen by mass than hydrogen. ● Does this mean that an oxygen atom is eight times the mass of a hydrogen atom? Atomic mass unit (amu) or (u): A unit based on the value of exactly 12 amu for the mass of the 12C isotope (6 protons and 6 neutrons in the nucleus). 1 amu = 1.66054 x 1024 g 1 g = 6.02214 x 1023 amu ● Since most elements have more than one isotope, a weighted average must be calculated to be representative of isotopic distributions found naturally. ● The average atomic mass of an element is determined by using the masses of its various isotopes and their relative natural abundances. An oxygen atom could have a mass of 16 amu, 17 amu , or 18 amu. Therefore a molecule of elemental oxygen, O2 , could have a mass of 32, 33, 34, 35, or 36 amu. To calculate Atomic Mass : 1. Multiply the decimal fraction of the natural abundance of each isotope by its mass (which is roughly equal to its mass number) 2. Add the results of these multiplications together Example: Carbon Natural distribution of carbon: 12C: 98.90% ** 14C is negligible 13C: 1.10% (0.9890 x 12.00 amu) + (0.0110 x 13.00 amu) = 12. 011 amu The atomic mass of carbon is 12.011 amu. ● Chemical symbol: a symbol assigned to an element based on the name of the element, consisting of one capital letter or a capital letter followed by a lowercase letter You must learn the first 36 element names and symbols, as well as any elements used in homework problems, class lectures, or lab. Properties of metals: 1. High thermal conductivity 2. High electrical conductivity 3. Malleability 4. Ductility 5. Has a metallic luster (Non Metals may occur as brittle, powdery solids or gases.) Atom versus an element: • An atom is unimaginably small. Most of an atom is empty space. The type of atom is determined by the number of protons, Z. • Elements can be individual atoms, like Ne. Some elements are diatomic (i.e. H2 ). • A sample of the element can have very different physical and chemical properties depending upon its structure, allotrope. Modern view of atomic structure ● The number of protons in the nucleus of an atom, called the atomic number, determines the identity of the atom and is characteristic for each element. ● For a neutral atom, the numbers of electrons and protons are equal. ● Ions have unequal numbers of protons and electrons. ● Atoms within an element (same atomic number) may have varying numbers of neutrons in the nucleus. ● While neutrons do not possess a charge, they have mass, therefore contribute to the overall mass of an atom. ● Atoms with identical atomic numbers (# of protons) but different mass numbers (different #’s of neutrons) are called isotopes. ● The sum of the protons plus the neutrons of an atom is called the mass number. ● Subtract the atomic number from the mass number to determine the number of neutrons of an isotope number. ● Isotopes may be represented symbolically by their elemental symbol, mass number, atomic number and when applicable, charge. ○ Isotopes can also be represented using the chemical symbol and the mass number: Ions and Ionic Compounds ● Ions: Ions are electrically charged atoms or groups of atoms called polyatomic ions. ○ Ions may be positively charged or negatively charged depending upon whether the atom has lost or gained electrons. ● Ions are represented by using the chemical symbol for the atom with a notation of the charge as a superscript to the upper right of the symbol. ○ Negative ions(called anions) have gained electrons ○ Positive ions(called cations) have lost electrons: Atomic Mass ● Since actual masses of elements are so small (i.e. 1.6735 x 1024g for 1H), it is easier to describe masses in atomic mass units. ● Atomic mass unit (amu) or (u): A unit based on the value of exactly 12 amu for the mass of the 12C isotope (6 protons and 6 neutrons in the nucleus). 1 amu= 1.66054 x 1024g 1 g = 6.02214 x 1023amu ● Since most elements have more than one isotope, a weighted average must be calculated to be representative of isotopic distributions found naturally. ● The average atomic mass of an element is determined by using the masses of its various isotopes and their relative natural abundances. An atom of the same element can have different mass properties. ○ An atom of C could have a mass of 12, 13, or 14 amu ○ The pure sample of C has the same chemical properties. ○ But do C12 and C14 have the same nuclear properties? ○ EX C12 is stable, C14 is unstable ● An oxygen atom could have a mass of 16 amu, 17 amu, or 18 amu. Therefore a molecule of elemental oxygen, O2, could have a mass of 32, 33, 34, 35, or 36 amu. To Calculate Atomic Mass: 1. Multiply the decimal fraction of the natural abundance of each isotope by its mass (which is roughly equal to its mass number) 2. Add the results of these multiplications together Elements and the Periodic Table ● Chemical symbol: a symbol assigned to an element based on the name of the element, consisting of one capital letter or a capital letter followed by a lowercase letter. ○ The names of elements, thus their chemical symbols, are derived from Latin, English, Greek, German, etc. names. ○ Elements have also been named in honor of scientists, or the geographic location where they were discovered. Avogadro’s number and the mole ● Chemists cannot directly “count out” the atoms of an element or molecules of a compound for chemical reactions. ○ The amu is impractical (too tiny) for laboratory use. ● In order to use large quantities of atoms or molecules and weigh out the amount of a substance needed, we can shift our thinking to the gram. ○ The mass of a single oxygen atom is 16 times the mass of a single hydrogen atom. ● Chemists use the mole (abbreviated mol) to relate the amount of a substance to the mass in “real world” quantities. ● Mole: The amount of substance that contains the same number of particles as there are atoms in 12 grams of 12C. ○ The mole is a collection of Avogadro’s number (6.022 x 10^23) of objects. ● Avogadro’s number: The number of atoms in exactly 12 grams of 12C to four significant figures: 6.022 x10^23. ● The lighter the atom, the less mass in one mol of atoms. ● Molar Mass: The mass in grams of one mole (mass in g/ 1 mol) of a substance.The molar mass of a compound is always numerically equal to its formula weight. The wave nature or light: (The tighter the waves, the more dangerous) The wave nature of light: ● The inverse relationship between wavelength and frequency is expressνλ = c ● c = speed of light (ms1) ● v = frequency (s1 or Hz) ● λ = wavelength (m) ● The shorter the wavelength, the higher the frequency ● The longer the wavelength, the lower the frequency ● ● The amplitude of the wave (maximum extent of the oscillation of the wave) relates to the intensity of the radiation. ● Electromagnetic waves have two components that travel in perpendicular planes (have the same wavelength and same speed): ● • Electric field component ● • Magnetic field component ● ● Example: A laser used in eye surgery to fuse detached retinas produces radiation with a wavelength of 640.0 nm. Calculate the frequency of this radiation. ● Use the relationship: c = n λ n = ? λ = 640.0 nm c = 3.00 x 108 m/s Note units: wavelength is often given in units of nm: 1 nm = 1 x 109 m or 1 x 109 nm = 1 m ● Three aspects of the behavior of electromagnetic radiation could not be explained by any one theory: a. 1. Emission of electromagnetic radiation from hot objects (blackbody radiation) b. 2. Emission of electrons from metal surfaces on which light shines (photoelectric effect) c. 3. Emission of light from electronically excited gas atoms (emission spectra or line spectra) ● Emission of light from hot objects blackbody radiation (molten metal, lava, etc) ● A black body is an idealized physical body that absorbs all incident electromagnetic radiation. ● A blackbody in thermal equilibrium emits electromagnetic radiation called blackbody radiation. At room temp. it appears black. However, if heated to a high temp. it will begin to glow with thermal radiation. ● Max Planck explained this phenomenon by postulating that energy can be released or absorbed by atoms only in discrete packets of energy (quanta). ● Quantum: The smallest increment of radiant energy that may be absorbed or emitted. ● Previously, it had been assumed that energy was continuous (transfer of any quantity of energy was possible). ● Planck asserted that energy could be gained or lost in wholenumber multiples of the quantity hn. ● Planck’s constant: h = 6.626 x 1034 Jouleseconds (Js) ● E = hn ● E = h c l ● Note that the energy of electromagnetic radiation depends upon the frequency: • The higher the frequency (shorter wavelength), the larger the energy value calculated (more damaging) Photoelectric Effect: the emission of electrons from a metal surface induced by light of a certain frequency. ● When light shines on a metal surface electrons are emitted – creating a current. ● • The light energy is transferred to the metal surface. Electrons are excited and gain kinetic energy. Electrons can gain so much energy they are ejected (ionized) from the atoms in metal. ● • There is a Threshold Frequency Below Which No Electrons Are ejected – no matter How Bright (intense)The light ● • If Light Were A Wave – Increasing the Intensity Should Increase The energy–And Eject electrons ● The next development in atomic structure was proposed by Albert Einstein, who used Planck’s quantum theory to develop an explanation of the photoelectric effect. Atomic Spectroscopy & Line Spectrum: The study of electromagnetic radiation absorbed and emitted by atoms. ● Sunlight, when passed through a prism, results in a continuous spectrum of colors. ● Atoms absorb energy and emit that energy as light. ● Atoms of each element emit light of a characteristic color. ● Each element has its own unique color combination (hydrogen above) (sort of like a fingerprint) ● Lasers produce light of only one wavelength (monochromatic light). ● Different types of lasers produce light of different wavelengths. i.e. different colors ● When voltages are applied across partially evacuated tubes containing gases, light of certain colors is emitted. Line spectra and the Bohr model ● Niels Bohr attempted to explain the line spectrum of hydrogen based on the idea that electrons move in circular orbits around the nucleus of an atom. ○ 1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies). ○ 2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom. ○ 3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = hn. ○ The lowest energy state is n=1 ○
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