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Week 3 Chemistry Notes

by: Amelia

Week 3 Chemistry Notes CHEM 1127Q-011


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Week 3 (9/12 and 9/14) Notes.
General Chemistry
Dr. Cady
Class Notes
General Chemistry, atoms and elements, Molecules and Compounds
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This 6 page Class Notes was uploaded by Amelia on Wednesday September 14, 2016. The Class Notes belongs to CHEM 1127Q-011 at University of Connecticut taught by Dr. Cady in Fall 2016. Since its upload, it has received 21 views. For similar materials see General Chemistry in Chemistry at University of Connecticut.


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Date Created: 09/14/16
9/12  CHAPTER 2­ Atoms and Elements (cont.)  The Atom (cont.)  1) Notation    a)  C­ The atomic symbol for the element  b) 6­ The atomic number, which represents the amount of protons in the element. The  atomic number is often left out of the notation because it is assumed the reader knows  this already.  c) 12­ The mass number shows the number of protons and neutrons, which also tells you  how much the element weighs.  d) The top right corner will display the charge of the element (Such as in O  )  2− e) The bottom right corner tells you how many of the atoms of that element there are    Properties of Subatomic Particles  Part of Atom  Charge  Relative Mass (amu)  Proton  +1  1.00728  Neutron  0  1.00867  Electron  ­1  0.00055    2) Isotopes  a) Isotopes are atoms with the same number of protons but different number of neutrons  i) For example, Carbon­12 consists of .9892(8) of all carbon, or 98.938%  b) Isotopes can be written with the symbol and atomic number or simply as the name  followed by a dash and the number  i) For example  can also be written as Neon­20  c) Most elements have isotopes   i) Hydrogen­1 makes up .999885(70) of all hydrogen  (1) Found in a typical water molecule  ii) Hydrogen­2 (Deuterium) makes up 0.000115(70) of all hydrogen  (1) Found in “heavy water”, which costs about $20 per 10g  iii) Hydrogen­3 (Tritium) is very rare  (1) It is radioactive and must be worked with in special labs  d) Atoms can have any number of isotopes  i) Example­ Calcium (Ca) has six different isotopes, and 96.941% is Carbon­40      3) Average Mass  a) Mass numbers are always whole numbers but the mass on the periodic table is a decimal  b) The mass on the periodic table is the average mass. You calculate it from the mass  number of all the isotopes and their natural abundance.  c) Atomic mass = (fraction of isotope 1 × mass of isotope 1) + (fraction of isotope 2 × mass of isotope 2) + …  i) If you have the percentage divide by 100 to get fraction.  d) Example Problem  i) Silicon has three natural isotopes  (1) Silicon­28 = 27.9769271amu makes up 92.2297%  (2) Silicon­29 = 28.9764949amu makes up 4.6832%  (3) The third isotope makes up 3.0872%  ii) Find the mass of the third isotope to four significant figures  (1) Atomic mass = (fraction of isotope 1 × mass of isotope 1) + (fraction of  isotope 2 ×mass of isotope 2) + (fraction of isotope 3 × mass of isotope  3)  (2) 28.09(amu of Si) = (27.9769271×0.922297) + (28.9764949×0.046832)  + (Si3× 0.030872)  (3) 28.09 = 25.8030399 + 1.3570272 + (Si3×0.030872)  (4) Si3 = 30.14amu, which is Silicon­30    Compounds, Ions, and the Periodic Table    1) Compounds (Molecules)  a) Molecular Formula vs. Empirical Formula  i) Molecular is the exact amount of atoms in a molecule   (1) Ex: C​ H26​​78​ 13 ii) Empirical is the simplest form of the formula (with whole numbers)   (1) Ex: C​ H26​​78​d13 ​e by 13 → C​ H​ O   2​ 6​ iii) Sometimes the molecular formula and the Empirical Formula are the same,  sometimes they are not  iv) Ex: Benzene’s molecular formula is C​ H​ , the6​m6​rical Formula is CH    2) Introduction to the Periodic Table   a) The rows across are called periods, and the columns down are called groups.   b) All elements want to be a noble gas (group 18) so they will want to lose or gain electrons  to have the same number as the nearest noble gas.  i) Group 17 (Halogens) like to have ­1 charge  ii) Group 1 (Alkali Metals) like to have +1 charge  iii) Group 2 (Alkaline Earth Metals) Like to have =2 charge  iv) (Noble gases are inert and will not interact with other elements)  c) This method of prediction becomes less accurate the farther you travel from the noble gas  or work with metals  d) Charge  i) Atoms with a positive charge are called Cations  ii) Atoms with Negative charge are called Anions  e) Some elements like to be in pairs, they form pure molecules.  i) H​2​O​2​F​2​Br​2​I2​ N2​ C2​    (1) An acronym to remember is HOFBrINCl   ii) These elements can usually be found in the Halogens group (group 17) or at the  top of group 16 and 15.  (1) Hydrogen is often written one top of the Halogens group as well as the  Alkali Metals because of it’s unique properties    3) Ions  a) When atoms or molecules lose or gain electrons they form charged particles called ions.  +​ ­ i) Na → Na​ +e​   ii) O+2e​ → O​   2­ b) There is no change to the protons in the element  c) Positive ions are called cations and negative ions are called anions    4) Molecules vs. Ions  a) Most molecules and atoms are happy with the number of electrons they have, so they do  not have a charge.  b) Ions are molecules/atoms that want to give up or take an extra electron or two  c) A sign that a compound is ionic is the first element in the molecule is from the first two  groups with a positive charge    5) Polyatomic Ions  a) “​A polyatomic ion, also known as a molecular ion, is a charged chemical species (ion)  composed of ​two​ or more atoms covalently bonded or of a metal complex that can be  considered to be acting as a single unit.’’ ­Wikipedia  b) Ammonium (NH​ )​ is the4​nly important positive polyatomic atom we need to remember  (though we should know hydronium (H​ O​ ).  +​ 3​ c) Hydroxide Ion (OH)​ is one of many negative polyatomic ions we need to remember. ​The  rest are in the book. Use Figure 2.5 (but if your book is published in 2015 use the online  version, your book is older.)      6) Ionic Compounds  a) Also known as an ​electrolyte ​(which is a sugar or salt)  b) Na​ Cl​ → table salt/sodium chloride is a strong electrolyte     Naming   1) Ions  a) Monatomic Cations  i) Basically the same word, unless there are different charges of that element, then  must designate with Roman Numeral  +​ (1) Na​  → Sodium  (2) K​  → Potassium  3+​ (3) Mn​  → Manganese (III)  b) Monatomic Anions  i) Basically add ­ide to the end of the word  ­ ​ (1) F​ → Fluoride  (2) H​ → Hydride  2­​ (3) O​  → Oxide  3­​ (4) N​  → Nitride    9/14  c) Oxoanions        ClO​  → perchlorate  4​ ­​   2­​ ­​ NO​  3​nitrate SO​ 4​ sulfate  ClO​  3​chlorate  ­​ 2­​ ­​ NO​  2​nitrite  SO​ 3​ sulfite  ClO​  2​chlorite      ClO​ → hypochlorite    i) ­ate = more oxygen  ii) ­ite = less oxygen  d) When going from name to formula ​always balance formulas  i) Example: Sodium Sulfide  (1) Na is ​  and S is ​   2­ (2) 2(+) + 2­ is zero → balanced  (3) Na​ S 2​ ii) Example: Iron(II) Oxide  2+​ 2­ (1) Fe is ​  and O​   (2) 2+ and 2­ are already balanced  (3) FeO  e) Summary  i) Cations: the name is the same as the element  ii) Anions: add ­ide to the name of the element   iii) Polyatomic: STUDY HARD ( ​ Flashcards?)  iv) When you a write a molecular formula for ionic compounds ​remember to balance  the charge.​   1) D  2) Binary Molecular Compounds  a) Unlike ionic compounds there is no single way to deduce the formula of a binary  molecular compound.  b) Systematic Naming  i) The first word is the name of the first element in the formula, with a Greek prefix  if necessary.  ii) The second word consists of:  (1) The appropriate Greek prefix  (2) The stem of the name of the second element  (3) The suffix ­ide  iii) Number of Prefixes  (1) 1 mono  (2) 2 di  (3) 3 tri  (4) 4 tetra  (5) 5 penta  (6) 6 hexa  (7) 7 hepta  (8) 8 octa  iv) Example  (1) P​ O​   4​6 (2) P is phosphorus, with the prefix tetra for 4  (3) O is oxygen, with the prefix hexa and the suffix ­ide  (4) Tetraphosphorus hexaoxide    3) Acids  a) Common Acids  i) You recognize an acid because it starts with hydrogen  ii) The names of acids are derived from the names of ions    Pure Substances (don’t use often)  Water Solutions (assume use this)  HCl(g) → Hydrogen chloride  H​ (aq), Cl​(aq) → Hydrochloric acid  +​ ­​ HBr(g) → Hydrogen bromide  H​ (aq), Br​(aq) → Hydrobromic acid  +​ ­​ HI(g) → Hydrogen iodine  H​ (aq), I​(aq) → Hydroiodic acid    iii) To get the name add Hydro and ­ic to the second element (the one that isn’t  hydrogen) and add the word acid  b) Oxoacids of Chlorine  i) Big ones to remember  (1) HNO​  → ni3​ic acid  (2) H​ SO​  → sulfuric acid  2​ 4​ ii) Basically every ­ate becomes ­ic and every ­ite becomes ­ous. The prefixes per­  and hypo­ stay the same    ClO​  → perchlorate ion  HClO​  → perchloric acid  4​ 4​ ClO​  3​chlorate ion  HClO​  → 3​loric acid  ­​ ClO​  2​chlorite ion  HClO​  → 2​lorous acid  ClO​ → hypochlorite ion  HClO → hypochlorous acid  (column one is the ion, column two is the acid)         


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