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Chemistry notes week 5

by: Annika Verburg

Chemistry notes week 5 Chem 113.01

Annika Verburg

GPA 3.8

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Class notes for unit 2
Introductory Chemistry
Dr. Aaron Robinson
Class Notes
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This 7 page Class Notes was uploaded by Annika Verburg on Monday September 26, 2016. The Class Notes belongs to Chem 113.01 at Abilene Christian University taught by Dr. Aaron Robinson in Fall 2016. Since its upload, it has received 3 views. For similar materials see Introductory Chemistry in Chemistry at Abilene Christian University.


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Date Created: 09/26/16
Unit 2 Chapter 4- Atoms and Elements Elements and Symbols  Now we turn to matter itself and we begin with the “building blocks” of matter: elements  There are 118 known elements – 88 naturally occurring and 30 man- made.  Each element is represented by a 1 or 2 letter symbol, as shown in the Periodic Table of Elements. Some symbols are obvious, others are not because they are based on Latin or Greek o One letter symbols are always capitalized:  Ex: carbon, C; iron, Fe; cobalt, Co; sodium, Na o Why is this important? CO is a combination of the elements carbon © and oxygen (O) and is a molecular compound called carbon monoxide. Co is the symbol for the transition metal element cobalt. The Periodic Table  The Periodic Table lists only elements – no compounds or mixtures  It is laid out in a very precise and information- packed way. o As we shall see, there are repeating (periodic) trends in the properties and behaviors of the elements o These trends were first noticed and published by the Russian chemist Dmitri Ivanovich Mendeleev in 1872  Periods and Groups o A horizontal row is called a Period. Periods are numbered from top to bottom  Period 1-hydrogen and helium  Period 2 – lithium and neon  Period 3- sodium through argon o A vertical column is called a group. Groups are numbered from left to right.  There are 3 numbering schemes: US, European, and IUPAC  We will mostly use the IUPAC system: Groups are number 1-18  The name group comes from the fact that all elements in a group behave similarly; similar chemical properties o Some groups have special names  Group 1 – Alkali Metals  Group 2 – Alkaline Earth Metals  Group 3 through 12- transition metals  Group 17 – halogens  Group 18 – noble gases o The Lanthanides and Actinides are special cases  Lanthanides are all in period 6 and group 3  Actinides are all Period 7 and Group 3 o Examples  What is the name and symbol of the element in period 4 and group 17? Br – bromine  What period and group does uranium belong to? Period 7 and group 3  What is the name of the period 3 halogen? Cl- Chlorine  In what period and group does the largest known Noble Gas belong? Period 7, noble gas  Metals, nonmetals, and metalloids o In addition to periods and groups, there are some regions of the Periodic Table which are important to know  The bold, zig-sag line running from B(element 5) to At (element 85) seperates the Metals on the left from the Nonmetals on the right Metals- most of the elements are metals (77%). They conduct heat and electricity well, have metallic luster, and lose electrons easily  Nonmetals- (15%) these are poor thermal and electrical conductors, usually have a dull, nonmetallic luster, and tend to gain electrons  Metalloids – (7%) these have properties intermediate between metals and nonmetals. o You should be able to identify an element as a Metal, nonmetal, or metalloid  Elements essential to health o Only 23 elements are essential for human health  Of these 4 major elements, 7 macrominerals, 12 micromenerals  Major elements – C, H, N, O  Macrominerals – Na, K, Mg, Ca, P, S, Cl  Microminerals – V, Cr, Mo, Mn, Fe, Co, Cu, Zn, Si, As, Se, I The Atom  Elements consist of one type of atom  Dalton’s atomic theory o All matter is made of tiny particles called atoms o All atoms of the same element are the same and those of different elements are different. o Atoms of two or more elements combine to form compounds. o A chemical reaction involves the rearrangement, separation, or combination of atoms in new ways. Atoms are never created or destroyed in a chemical reaction  Electrical charges and subatomic particles o Atoms are made of protons and neutrons o Nucleus discovered by Ernest Rutherford  In an atom, the protons = electrons, so that the positive charges = negative charges  These masses are so small that an appropriate sized mass unit was invented, called the Atomic 24ss unit o 1 amu = 1.6605 * 10 grams Atomic number and mass number  atomic number – is the number of protons in an atom o the #p determines the identity of an atom o the number of neutrons in an atom may vary, but this does not change its identity.  Mass number – an atom’s mass number is the sum of protons and neutrons in the atom. (mass # = p+n) Isotopes  Isotopes – atoms with the same element but with different number of neutrons o They have the same Atomic number, but different Mass Numbers  Atomic Mass – weighted average of the masses of all isotopes of an element o Weighted means that each isotope counts or is “weighted” in the average by its abundance Electron Energy Levels (Electron Arrangement in Atoms)  We now turn to the way in which the electrons in an atom are arranged or “structured”: electronic structure.  Electron Energy Levels – Electrons have very specific energies which depend on the type of atom and the distance of the electron from the nucleus. These are called Energy Levels. o Energy Levels – These are also called Principal Energy Levels or Shells  The Levels are numbered n = 1 to n = 7 as one moves outward from the nucleus  These levels increase in energy as you move further from the nucleus. Electrons in the first shell (n = 1) are lowest in energy; electrons in n = 2 are higher energy  The energy of a level is very precise.  EX: the energy of electrons in the n = 2 level of phosphorus are the same in any P atom, and are different from the electrons in the n = 2 level of carbon or sulfur  Changes in Electron Energy Level o To “jump from one energy to another electron must gain or lose energy o Usually this happens due to an atom absorbing or emitting light  Absorb light energy => move to higher energy level  Emit light energy => move to lower energy level  Color is due to this process  Sublevels o Each energy level contains one or more sublevels o There are 4 types of sublevels denoted s, p, d, and f o The number of sublevels in each energy level  Level 1: 1 sublevel (s)  Level 2: 2 sublevels (s,p)  Level 3: 3 sublevels (s, p, d)  Level 4: 4 sublevels (s, p, d, f) o Each sublevel contains one or more orbitals  S sublevels: one s orbital  P sublevels: three p orbitals  D sublevels: five d orbitals  F sublevels: seven f orbitals  Orbitals o Orbitals are a sort of volume which holds electrons. More precisely, they are probalitlity volmues within whih there is a high porbablilti for an electron for an electron to reside o There are 4 types of orbitals, each with a unique shape. They are names like the sublevels,  S orbitals are spherical and exist in the s Sublevels  There is one s orbital in an s Sublevel  S orbitals always occur as “singles” (in sets of one)  Each energy level contains a single s orbital o 1s is the name of the s orbital in the n = 1 level; 2s in the n = 2 level, etc. o orbitals get larger as the level increases  p orbitals are “dumbbell- shaped’ and exist in the p sublevels  there are three p orbitals in a p sublevel  p orbitals always occur as “triplets” (in sets of three)  energy lever n = 2 and above each contain three p orbitals o 2p is the name of the p orbitals in the n = 2 level; 3p in the n = 3 level et. o Orbitals get larger as the level increases  D orbitals are “X-shaped” and exist in the d sublelvels o There are five d orbitals in a d sublevel o D orbitals always occur in sets of five o Energy level n = 3 and above each contain five d orbitals  3d is the name of the d orbitals in the n =3 level; 4d in the n = 4 level  orbitals get larger as the level increases o f orbitals are “8-lobed” and exist in the f sublevels  there are seven f orbitals in an f sublevel  f orbitals always occur in sets of seven  energy level n = 4 and above each contain seven f orbitals  4f is the name of the f orbitals in the n = 4 level; 5f in the n = 5 level o orbitals get larger as the level increases  Orbital capacity and Number of Electrons in Sublevels o Each orbital can hold a maximum of 2 electrons  A orbital may contain 0, 1, or 2 electrons, but never more than 2  An orbital with 0 electrons in vacant or empty  An orbital with 1 electron is half-filled  An orbital with 2 electrons is filled  Thie is true of s, p, d, or f: each can hold 0, 1, or 2 electrons  A sublevel can hold a maximum of 2 electrons in each of its orbitals o S sublevels can hold 2 electrons max o P sublevels can hold 6 electrons max o D sublevels can hold____ electrons max (in its set of ______ d orbitals) o F sublevels can hold_____ electrons max (in its set of ______f orbitals)  How many electrons can each energy level hold, max? o The n = 1 level o The n = 2 level o The n = 3 level o The n = 4 level Electron Configuration  Electron configurations are ways of showing how electrons are arranged or distributed within an atom. Another, related way is with orbital diagrams. o Orbital diagrams give slightly more information, but o Electron configurations are more compact.  Orbital diagrams – boxes or lines represent orbitals. Arrows represent electrons o Electrons can be in an orbital either “spin up” or “spin down”  Thus there are 2 possible ways or putting electrons into an orbital, and each is unique  That is why an orbital can only contain 2 electrons max: one spin up and spin down  Sublevels fill with electrons in order of increasing energy: lowest energy sublevels fill first  There is a specific order of energies  The “filling order” is shown in the Periodic Table o Filling Rules:  Lowest energy first  Opposite spins for electrons in the same orbital  Same spin for electrons in different orbitals in a sublevel o Blocks  Four regions of the Periodic Table: s block, p block, d block, f block  The s block o The first two columns (group 1 &2) o Represent filling the s sublevel in each energy level  The p block o The last 6 columns (group 13 – 18) o Represent filling the p sublevel in each energy level  The d block o The transition metal columns (group 3 -12) o Represent filling the d sublevel in each energy level  The f block o The columns below the main table  Writing Electron Configurations using sublevel blocks o We can use the periodic table to help write electron configurations  The PT is arranged in filling order, so there is no need to memorize  By learning to read the s, p, d, f blocks in the PT, we can simplify the writing of the Electron Configurations and can provide ourselves with a “double check method


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