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CHEM 100 CH 3 notes

by: Carly Holliday

CHEM 100 CH 3 notes Chem 100

Marketplace > Indiana State University > Chem 100 > CHEM 100 CH 3 notes
Carly Holliday
GPA 3.3

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All of the vocab and challenging things we learned in CH 3 are all right here! Happy Studying!
Chemistry 100
Dr. Jeewandara
Class Notes
Chemistry, Moles, valence electrons, ions, Isotopes
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This 6 page Class Notes was uploaded by Carly Holliday on Sunday February 14, 2016. The Class Notes belongs to Chem 100 at Indiana State University taught by Dr. Jeewandara in Summer 2015. Since its upload, it has received 19 views.


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Date Created: 02/14/16
CHEMISTRY IN FOCUS CHEM 100 CHAPTER 3 NOTES 3.1 A Walk on the Beach Think of the things you smell and hear and see and feel on a beach. All of those things are made  up of atoms. Atoms are all around us. They are the building blocks of our life and our world.  To understand the world we see and how it connects to the world we cannot see (the microscopic world) we have to understand atoms and how they function. There are 91 types of atoms. They are arranged on the periodic table of elements. There are some atoms that do not react with anything to make a compound. They are found in  nature as pure substances. Other atoms react with almost everything and make up all of the  compounds we have. They are not found alone in nature. Mass vs. weight:  Mass: the measure of quantity of matter (the same everywhere in the universe) Weight: how strong gravity pulls on an object (different on earth and on the moon) **Things to remember** 1. All samples of a compound have the same proportions of their constituent elements. 2. All matter is made of atoms  3. Atoms differ from one element to another  3.2 Protons Determine the Element The number of protons in the nucleus determines what the element is. There are three subatomic particles in an atom  Protons  Neutrons  Electrons Protons (+): Positively charged and make up part of the nucleus.  Neutrons (0): Neutrally charged and are also found in the nucleus. Electrons (­): negatively charged and found in the area surrounding the nucleus.  Charge: Positive or negative. These opposite forces act on one another to keep all the particles  together. Nucleus: Very dense center of the atom. It is made up of protons and neutrons and carries most  of the mass of the atom. It is positively charged.  Atomic Number: This number is always the number of protons in an atom. It is on the periodic  table above each element. Chemical Symbol: This is the abbreviated letter(s) on the periodic table that stand for the  element it is assigned to. It is usually made up of one CAPITAL letter by itself or a CAPITAL  lowercase combination. The second letter is always the lowercase one. Amu: Atomic mass unit. The atomic masses found on the periodic table are the weight averages  of the naturally occurring isotopes for an element.  3.3 Electrons A neutral atom has as many electrons outside of its nucleus as protons inside of its nucleus. An atom can, however, lose or gain one or more of its electrons. When this happens, the charges  of the electrons no longer exactly cancel the charges of the protons, and the atom becomes a  charged particle called an ion. Cation: A positively Charged Ion for example: Na + Anion: A negatively charged Ion for example F ­ Electron configuration: Which orbits the electrons fill.  3.4 Neutrons Neutron: A subatomic particle found in the nucleus. It has a neutral charge (0). Sometimes atoms of the same substance have the same number of protons but different numbers  of Neutrons. When this happens the atom is called an Isotope. An Isotope and a regular atom of the same substance have different masses because there are  extra neutron(s) in the nucleus.  Mass Number: The sum of protons and neutrons in an atom.  When there are more neutrons, the mass number is higher.  Example:    1C  is an isotope of Carbon. The 13 is the mass number.  We can find the number of protons from the periodic table p = 6 then we can subtract 6 from  13 13 – 6 = 7 Neutrons.  13 In Isotope  C  there are 6 protons and 7 neutrons. 3.5 Specifying an Atom A C  Here is a visual    X A the Mass Number (Protons + Neutrons) Z the Atomic Number (number of protons) C the overall Charge (if the atom has more or less electrons than protons) X the Chemical Symbol Example:  If  A=17 and Z= 9  and C=0 Protons 9  Electrons 9 Mass #  17 Neutrons 17 – 9 = 8 3.6 Atomic Mass How to calculate Isotope1 %isotope2 %Isotope3 Average Atomicmass= 100 xmassIsotope1+ 100 xmassisotope2+ 100 xmass Isotope3 Example: Magnesium has three naturally occurring Isotopes  Masses: 23.99 amu, 24.99 amu, 25.98 amu Natural abundance: 78.99%, 10.00%, 11.01% 78.99 10.00 11.01 Calculation 100  x 23.99 +100  x 24.99 +100  x 25.98 = 24.309099amu 3.7 Periodic Law Dmitri Mendeleev (1834–1907):  He noticed that some elements had similar properties, and he  grouped these together.  Example 1: helium, neon, and argon are all chemically stable gases and could be put into one  group.  Example 2: Sodium and potassium are reactive metals and could be put into another group. Mendeleev did not know why the periodic law existed. His law, like all scientific laws,  summarized a large number of observations but did not give the underlying reasons for the  observed behavior. The next step, following the scientific method, was to devise a theory that  explained the law and provided a model for atoms. 3.8 Explanation of the Periodic Law Bohr Model: # of protons Identity  # of electrons  Chemical Behavior  Electrons orbit the nucleus like the planets orbit the sun Bohr specified each orbit with an integer n, called the orbit’s quantum number. The higher the  quantum number, the greater the distance between the electron and the nucleus and the higher the electron’s energy. Each orbit can only hold a certain number of electrons. Orbit 1 (n=1) can only hold 2 electrons  Orbit 2 (n=2) can only hold 8 electrons  Orbit 3 (n=3) can only hold 8 electrons Valence electrons: the electrons found on the outermost orbit. 3.10 Families of Elements All of the Groups are colored on the table below: Group 1 (IA) has one valence electron Group 2 (IIA) has two valence electrons Group 3 (IIIA) has three valence electrons Group 4 (IVA) has four valence electrons Group 5 (VA) have five valence electrons  Group 6 (VIA) has six valence electrons Group 7 (VIIA) has seven valence electrons Because they have the same number of valence electrons, they have similar reactivity qualities. These Groups have special names: Alkali Metals: Group 1 Alkali earth Metals: Group 2 Chalcogens: Group 6 Halogens: Group 7 Noble Gasses: Group 8 Diatomic Molecules: 1 dozen = 12 2 dozen = 24 1 Mole = 6.022x10 23 2 Moles = 12.44x10 23 23 ½ Mole = 3.11x10


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