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Regents Chemistry Mrs. Ingersoll Unit 1 – Introduction to Matter and Atomic Structure 2 Matter has mass....... .....and volume can be divided into substances + mixtures Substance matter with constant properties and composition, regardless of its source – it is the same throughout can be divided into elements and compounds 3 Element A substance that can not be broken down to simpler substances Made of only one kind of atom Examples: gold carbon oxygen 4 Compounds: --have two or more kinds of atoms held together by chemical bonds Examples water carbon dioxide sodium chloride glucose H2O CO2 NaCl C6H12O6 -- formed by a chemical change -- properties of the original elements are no longer apparent -- can be decomposed by a chemical reaction 5 Mixtures --variable properties and composition can be: homogeneous Or heterogeneous -- evenly mixed -- not evenly mixed Salt water, Kool-aid Any solution Chocolate chip cookies, Raisin bran --form without a chemical change --properties of ingredients still apparent --can be separated without a chemical reaction by using the physical properties of the substances in the mixture 6 Mixtures: Separation Techniques 1. Filtration – based on differences in particle size, or solubility 2. Distillation – based on differences in boiling point 7 3. Chromatography – based on differences in solubility and attraction to another substance, ex. paper Differences in density, freezing point, and other properties can also be used to separate mixtures 8 SUMMARY Element Substances Compound Mixture element + element element + compound compound + compound 9 THE STRUCTURE OF THE ATOM Democritus, 400 BC --Matter is made of small, indivisible particles – “atomos” In the 1700’s, scientists making accurate measurements discovered several new laws -- 1. Law of Conservation of Matter --matter is not created or destroyed during chemical or physical changes 2. Law of Definite Composition -- compounds have an unvarying composition 3. Law of Multiple Proportions -- elements combine in simple ratios 11 John Dalton, 1807 -- Atomic Theory 1. Each element is composed of one kind of atom 2. Atoms of different elements have different masses and properties 3. Atoms can combine and recombine with each other; but they don’t change from one kind to another; they aren’t created or destroyed 4. In compounds, atoms combine in simple, constant ratios Dalton thought that atoms were hard, unbreakable spheres 12 J.J. Thomson, 1897 --experimented with cathode ray tubes: vacuum tubes with high voltage passing through --found that cathode rays could be: 1. deflected by magnets 2. bent toward positive electric fields CONCLUSIONS: 1. cathode rays were streams of negatively charged particles electrons 2. atoms are made of sub-atomic particles with opposite electrical charge 13 14 Ernest Rutherford, 1910: Gold Foil Experiment -- took alpha particles (2 protons + 2 neutrons) and shot them toward a thin gold foil 15 Rutherford’s Results: Most of the particles went straight through; a few were deflected 16 Interpretation of Rutherford’s ExperimentHis observations and conclusions: 1. Since most of the particles went straight through, the atom is mostly empty space. 2. Some of the alpha particles were deflected so the atom contains a nucleus that is: a. very small b. dense (a large mass in a small volume) c. positively charged 17 The nucleus contains most of the atom’s mass and all its positive charge proton mass 1u neutron 1u 0 Atomic number Atomic mass (mass number) charge +1 = number of protons determines the element = number of protons + neutrons atomic mass atomic number 10.81 B 5 18 Not all atoms of the same element are identical p p n Atomic number: 1 Atomic mass: 1 H-1 n p n 1 1 2 H-2 3 H-3 These are called isotopes of hydrogen Isotopes: atoms of the same element with different masses -- same p, different n 19 Isotopes of hydrogen have the same chemical properties All are: -- flammable gases --able to react with oxygen to form water They differ in mass and nuclear properties All elements exist as mixtures of isotopes -the atomic mass of an element is the mass of an average atom A carbon-12 atom has a mass of 12.0000000... yet the atomic mass of the element carbon is 12.011 because a few carbon-13 and carbon-14 atoms are mixed in 20 Sample Problem Element X consists of 35% X-21 and 65% X-23 What is the average atomic mass of X? Solution Assume you have 100 atoms: 35 of them would be X-21 65 of them would be X-23 Find the total mass of all 100 atoms: 35 x 21 = 735 65 x 23 = 1495 2230 u = mass of 100 atoms 2230 / 100 = 22.30 u The average is always between the isotope masses and always closer to the one that is most abundant 21 Charged Atoms + charge in nucleus, – charge outside Add + and – to find the total charge If the charge = 0 ( p = e ), it is a neutral atom If the charge = 0 ( p = e ), it is an ion Describing Atoms: 1. Atomic Number = number of protons 2. Atomic Mass (or mass number) = number of protons + neutrons 3. Electrical Charge = protons(+) + electrons(-) a neutral atom has a charge of 0 (protons = electrons) a positive ion (+ charge) has more protons than electrons a negative ion (- charge) has more electrons than protons 22 ee- Examples: atomic number = _________ 3 atomic mass = ________ 7 = proton charge = _______ +1 = neutron element = ______ Li 10 How many neutrons are in an atom of F? _______ 5 How many electrons are in a neutral atom of boron? ______ Suppose an atom has 14 protons, 16 neutrons, and 16 electrons Si -- what element is it? __________ 14 -- atomic number = _______ -- atomic mass = ______ 30 -2 -- charge = _______ 23 HOW TO READ THE PERIODIC TABLE OF THE ELEMENTS 24 ELECTRONS Niels Bohr, 1913 -- said that electrons orbit the nucleus in set paths, like planets around the sun: or These orbits were designated: K L M N O closest to nucleus lowest potential energy P Q farthest from nucleus highest potential energy 25 Bohr said that e- could be in any of these 7 orbits, but NOT in between. Add energy e- move to higher energy levels, absorbing the energy Later, the e- falls back to a lower energy level, and the energy is released in the form of electromagnetic radiation e e e energy absorbed energy released 26 Electromagnetic Radiation -- energy that travels as waves -- can travel through a vacuum -- moves at 3.0 x 108 m/s, or “c” (speed of light) λ -- wavelength -- the number of waves that move past a fixed point in one second is called the frequency, “f” -- frequency is measured in Hertz (Hz), or waves per second 27 -- As λ decreases, f increases -- c = λ f 28 29 In 1900, Max Planck discovered: the energy carried by electromagnetic energy is directly proportional to its frequency (high frequency means high energy) E= h f energy = (planck’s constant)(frequency) --so, as λ , f , E 30 Back to atoms .... When all the e- in an atom are in the lowest possible energy orbits, the atom is in the ground state – its normal condition Add energy, and the e- move to higher energy levels, and the atom is in the excited state – an unstable condition Energy is released when the e- return to the ground state L K jump--- short, little energy released light emitted has low f, high λ red M K L M N N K jump--- medium energy released light emitted has med f, med λ green K jump--- long, lots of energy released light emitted has high f, low λ violet 31 Each electron jump releases one photon of a certain λ This produces the spectrum of the element Since each element starts with a different electron arrangement, each element has a different spectrum 32 Bohr’s model worked well for hydrogen, but not for more complex elements. Other scientists in the 1920’s came up with several modifications to his theory. The Current Understanding of Electron Arrangement: The Wave-Mechanical model We can’t know an electron’s distance from the nucleus, as Bohr thought, but we can calculate how likely it is to be found at certain distances K L electron probability electron probability probability probability probability distance from nucleus distance from nucleus distance from nucleus 33 The wave-mechanical model says that electrons are in an electron cloud – there are NO orbits Instead, we have Principal Energy Levels or Shells: K 1 L 2 M 3 N 4 O 5 P 6 Q 7 Each level can hold a given number of electrons: PEL No. of e1 2 2 8 3 18 4 or higher 32 2 If the PEL = n, the number of electrons = 2n 34 The Principal Energy Levels are divided into orbitals (NOT ORBITS) of different energy An orbital is the region where the electron is most likely to be found 35 Rules for electrons in atoms: 1. No 2 electrons in an atom have exactly the same energy 2. Octet Rule – the last shell can never hold more than 8 e- Example: Atomic No. 18 Ar: 2 – 8 – 8 Atomic No. 19 K: Atomic No. 19 K: 2–8–9 2–8–8–1 3rd shell can hold 10 more e violates octet rule 4th shell begins to fill So, a shell may not be filled with electrons before the next shell begins to be filled 36 - 15.9994 Electron configurations on the periodic table are for: Neutral atoms-- Ex. oxygen is 2–6 -2 O 8 2–6 when its charge is –2, the electrons are… 2–8 In the ground state -- Ex. oxygen is 2–6 oxygen is 2—5—1 in the excited state 37 Atoms are most stable when their outermost energy level has as many e- as possible (8 e- ; or 2 e- for the 1st shell) This situation is called a stable octet Elements that have a stable octet when neutral are called the noble gases (group 18) These elements undergo few, if any, chemical reactions Other elements change during chemical reactions to obtain a stable octet 38 Electrons in the outermost energy level are called valence electrons Ex. Atomic No. 13 Al 2 – 8 – 3 Lewis electron-dot structures: -- use one dot for each valence electron Ex. Al Ne Li 39