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by: Esraa Hagag

chemistry chem135

Esraa Hagag
Northampton Community College

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these will cover week 4 and 5
chemistry of life
edward s fleming
Class Notes
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This 8 page Class Notes was uploaded by Esraa Hagag on Thursday September 22, 2016. The Class Notes belongs to chem135 at Northampton Community College taught by edward s fleming in Fall 2016. Since its upload, it has received 4 views. For similar materials see chemistry of life in Chemistry at Northampton Community College.


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Date Created: 09/22/16
Separation of mixtures: A mixture can be separated into pure substance components on the basis of different  component physical properties 1. Distillation  Separation of liquid mixture based on different component boiling  points.    e.g. distillation of liquor, petroleum refining 2. Crystallization  Separation of soluble substance from a solution based on component  solubility by controlling temperature and the amount of solvent, we can cause a soluble substance to crystalize or fall out of solution.            e.g. “rock candy” from sugar solution 3. Filtration  Remove insoluble substance from a liquid Structure of the atom: All atoms contain three sub atomic particles  Proton P  Neutron N  Electron e Mass: is in atomic mass unit (amu)=1.66*10  g­24 Charges: like charges repel and unlike charges attract each other Mass number (A) A= number of protons + number of neutron Atomic number (Z) Z= number of protons Z is unique for each element  e.g. Z=1 ­­­­­­­­­­­­ hydrogen  In neutral atom number of protons= number of electron  If number of electron less than number of protons we have a positively charged  particles  If number of electrons greater than number of protons we have negatively  charged particles Isotopes  Number of neutrons in a given element tells us what isotope of element we have  For a given element the number of protons is always same, but the number of  neutrons is variable  Isotopes of a given element have the same chemical and physical properties Generic representation for isotope  Look at page 111 in your book Atomic weight:  It is a weighted average of its isotopes  Elements are found in nature as a mixture of their isotopes in fixed ration  Mass of any isotope is less than its mass number Discrepancy: is due to “mass defect” Mass defect: means that some mass is lost by conversion to energy Energy: binding energy which holds the nucleus together Periodic Table It is a classification of element based on two things 1. Increasing atomic number 2. Repeatability of properties or property trends Periodic table: 1. It has 18 columns or group up and down  Atomic size increases as we go down particular group  Elements in a particular group often have similar properties  E.g. group 1A (except H) are all alkali metals  They all have low melting points  They all soft  They all react with water  They form ionic compounds with non­metals  E.g. group 7A Halogens  Relatively strong oxidizing agents  Form ionic compounds with metals  E.g. group 8A Noble (Inert) gases  Unreactive chemically  Often used in fluorescent lights (krypton, neon, xenon) 2. It has 7 rows (periods)  Properties vary greatly across a period  A cross a period (left to right) atomic size decreases  The added electrons are in the same shell and pulled closer by proton’s in  nucleus Main group elements Transition elements (metals) Inner transition elements  Groups 1A and 2A (s   B group elements (d   At bottom of periodic  block) block) it separates  table to keep it from   Groups 3A­ 8A (p  main group being too wide block) Classification of elements: Metals Non metals metalloids  Conduct   Do not conduct   Properties between metals  electricity electricity except  and no metals  Shiny or would  carbon in the form of   Use as semi­conductor be if scraped graphite e.g. silicone (si) +  germanium (ge) Electronic structure of atoms:  Planetary model is inadequate to explain observed behavior of electrons  New model needs to account for  As electrons get further away from nucleus they have more energy  Allowed electrons energies are not continues  Electrons energies are discrete, quantized, separate  Electrons are therefore allowed in only certain energy levels  Electrons can change energy levels by absorbing energy or releasing energy  Absorbing energy: heated, absorbing certain light, by being discharged with  electricity  Energy levels are given by numbers Orbitals:  Locations of electrons associated with certain energy levels  Locations are not fixed but “clouds” or “envelopes” in which high probability of  finding electrons  Not all orbitals have same shape  E.g. s  orbitals­­­­­­­ spherical, p  orbitals­­­­­­­ dumb­ ben shaped  D & f orbitals­­­­­­­­­­ complex shape  At room temperature, an atom’s electrons are all in lowest energy levels (ground  state)  When electrons energy they go to higher energy levels (excited state) Electron configurations:  Complete list of all orbitals occupied by atom’s electrons in ground states  Number of protons= number of electrons  As Z increases number of electrons increase  Added electrons are further away from nucleus  End up with inner electrons and outer electrons  Outer electrons: individually in chemical reactions and responsible for some physical properties  Periodic table is arranged so that elements in a given group have some number of  electrons  For main group elements the group number give the number of valence electrons Ionization Energy (IE)  It is a periodic property which means it has trans with periodic table specifically  down a group IE goes down  If atom loses an electron we get a positive ion  Energy required to remove an electron is ionization energy from an atom in gas  phase  Across a period from left to right IE increases  Electron dot symbols: are used to defect an atom explicitly showing its valence  electrons  Look at pages 121, 122, and 123 figure 4.13, 4.14 Chapter 6 Compounds and Bonding Compounds:  Pure substances, when atoms combine to form compounds they are held  together by chemical bonds Chemical bonding:  Occurs when valence electrons are either transformed or shared between atoms Ionic bonding:  When valence electrons are fully transformed from one atom to another  Results in formation of ions  Cation: Element lose electrons positive ions  Anion: Element gains electrons negative ions  E.g. NaCl sodium chloride + Na   gives up one electron to form Na  which is sodium ion Cl    gains one electron to form Cl  ion which is chloride Ions are held together by electrostatic forces Naming monatomic ions: Cations:  Name of element and add the word “ion” + e.g. k  potassium ion, the + is the charge number  some elements can have more than one charge number e.g. Fe  iron (2) ion ­­­­­ ferrous ion Anions:  take name of element but use “­­­­ide” suffix 2­  O  oxide ion the (2­) means it gained 2 electrons Ionic compounds:  Consist of ­ & + ions  Overall charge is neutral (number of + charges is equal to number of – charges)  which means it is electrically neutral  Consist of 3D average of ions arranged in a repeating geometrical pattern (often  called crystal lattice or crystalline structure)  Look at page 171 in your book   Macro scale—bulk matter  Micro scale—optical magnification  Nano scale—scale on which ions, atoms, and molecules live  Since ionic compounds exist as crystal lattices, number molecules of ionic  compounds  Formula Nacl is not molecular formula  Formula unit: gives lowest whole number ratio of ions in crystals  Sometimes it is helpful to think of a formula unit represent a kind of molecule Why does atoms bond?  The most stable of un reactive element are in group 8A (8 valence electrons  corresponds to a complete outer shell)  Main group elements lose, gain or share electrons to get a complete outer shell  (octet rule)  Except period one which is hydrogen and helium, they need only two electrons to complete the outer shell  Look at page 175 in your book figure 6.2  In ionic compounds metals form positive ions and nonmetals form negative ions  Look at page 169 in your book table 6.2  Table summary:  Groups 1A, 2A, 3A, lose 1e, 2e, 3e   respectively  To get octet having the same outer shell configuration as preceding  noble gas  Groups 5A, 6A, 7A gain electrons 3e, 2e, 1e respectively  To get the outer shell configuration of next noble gas  Group 4A is a “grey” area  Octet rule is only a guide line  Applies very well to main group elements  Does not work at all with transition metals  Look at page 171 in your book number 6.1, 6.2 Predicting ionic compound formulas:  Use group number and octet rule to predict charge number  Use charge neutrality to predict the number of each ion Polyatomic ions:  Ions consisting of more than one atom  Look at page 178 in your book table 6.7  Hydroxide, ammonium, nitrate, carbonate, hydrogen carbonate, cynide,  acetate, sulfate, and phosphate Page 182 number 6.34 Covalent bond:   Atoms share one or more pair of valence electrons to get octet  Electrons are not transferred  Covalently bonded compounds exist as molecules General rules for prediction compounds bonding: 1. Metals+ nonmetals = ionic bond 2. Nonmetals +nonmetals = covalent bonds 3. Nonmetals + metalloids = covalent bonds Nonmetal diatomic molecules:  Many nonmetals exist as diatomic molecules   e.g. O 2 multiple covalent bonds:  most covalent bonds are single  C, O, N can form multiple  Covalent bond with each other P, S, Se can form multiple bonds with carbon,  oxygen, and nitrogen but not among each other   Covalent bonds only mean bonding taking place among nonmetals  Look at page 183 in your book number 6.10 and table 6.11 Lewis or electron dot structure:  Show all atoms  Show all bonding electrons  Show all nonbonding or long pair electrons Naming covalent compounds (binary compounds):  Name first element  Use prefix to indicate the number of atoms of this first element except mono  not used  Table 6.12 page 186  Then name the second element  Use “­­­ide” suffix  Use prefix to show number of atoms except mono not used (except carbon  monoxide)  Page 188 number 6.42, 6.46 Electro negativity:  In covalent bonds the electrons shared between atoms  When two atoms same electrons shared equally e.g. O , F ,2  2 2  When two atoms are different then bonding electrons are not shared equally  Electrons are closer to one of the atoms  Bonding electrons are closer to F since F is more electronegative than H  Electronegativity is a measure of an atoms ability to attract bonding electrons  to itself in a covalent bond  Page 189  Least is CS which electronegativity is 0.7 and most is ferrous which has  electronegativity 4.0  Electronegativity is periodic property  E.g. across a row in periodic table from left to right the electronegativity  increase Bond polarity:   Differences in electronegativity between two bonded atoms lead to unequal  sharing of electrons  Results in a no uniform charge distribution  If two bonded atoms are same different electronegativity = 0  Bond is nonpolar  If two bonded atoms have a sufficiently different electronegativity, then results in  a polar bond  Magnitude in difference of electronegativity relates to how strongly polar bond is


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