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CH 105 Chapter 4-6 Study Guide and Practice Problems

by: kmbaars

CH 105 Chapter 4-6 Study Guide and Practice Problems CHEM 105

Marketplace > University of Alabama at Birmingham > Science > CHEM 105 > CH 105 Chapter 4 6 Study Guide and Practice Problems
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Included are the main topics covered in chapters 4-6 as well as 19 math practice problems
Introduction to Chemistry I
Steven Harville
Study Guide
chemisty, bonds, electronegativity, atomic orbitals, atomic orbital shapes, Subatomic Particles, Isotopes, electromagnetic, spectrum, lewis dot, Ionization energy, atomic radius, ionic, compounds, covalent, ions, VSEPR, VSEPR shapes, polar, nonpolar
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This 12 page Study Guide was uploaded by kmbaars on Tuesday October 11, 2016. The Study Guide belongs to CHEM 105 at University of Alabama at Birmingham taught by Steven Harville in Fall 2016. Since its upload, it has received 30 views. For similar materials see Introduction to Chemistry I in Science at University of Alabama at Birmingham.

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Date Created: 10/11/16
CH 4-6 Study Guide  Dalton’s Atomic Theory: o Each element is made up of individual particles called atoms, which are indivisible- they cannot be created nor destroyed o All atoms of each element are identical, and atoms of one element are different from other elements o Compounds are formed by 2 or more atoms of different elements  Other atomic models: o Thompson Plum Pudding Model: negative electrons are surrounded by a “positive pudding” o Rutherford’s Atomic Model (Correct Model):  Atom composed of a dense nucleus and electron cloud; Positive charge concentrated in a very small space  Nucleus contains >99.99% of mass, but almost none of the volume  Characteristics of subatomic particles o Each atom has a fixed number of protons, called the atomic number of the element o Neutrons may vary between ions o When atoms lose/gain electrons, they become an ion o Each element is electrically neutral (has the same number of protons and electrons) Particle Symbol Charge Mass Location Proton p+ +1 1 Nucleus Neutron n 0 1 Nucleus Electron e- -1 .0005 Outside in electron orbitals o Number of Protons+ Neutrons= mass number  Isotopes: atoms of the same element with different mass numbers (different number of neutrons). The mass number, NOT the atomic number, will vary from isotope to isotope. o To calculate average atomic mass, multiply the experimental percent abundance by the atomic mass of that isotope, then sum the total mass of all isotopes. o Take the percent of relative abundance and convert it into a decimal by moving the decimal two places to the left o Multiply the new decimal by the mass that it goes with o Record the number and repeat the above steps for all isotopes given by the problem o Add all values for isotopes together to find the average atomic mass  Electromagnetic Spectrum: Our eyes can only see light in the minute visible light spectrum with medium energy wavelengths. A wavelength determines the amount of energy and is the length from crest to crest, or trough to trough on the spectrum.  Electron energy levels: as energy level increases, energy levels are closer together. Electrons absorb energy to jump to a higher energy level, which have sublevels o S Orbital: sphere shaped nucleus; single orbital in each s sublevel which can only hold 2 electrons o P Orbital: two-lobed shape; 3 p orbitals make up each sublevel which can hold up to 6 electrons o D Orbital: four-lobed and one two-lobed with a donut ring around it shaped; each d sublevel contains 5 sublevels which can hold up to 10 electrons o Sublevel Filling: Electrons fill orbitals from lowest to highest energy with a maximum of two electrons in each orbital. This is called electron configuration  Lewis dot structures represent the number of valence electrons in an element o Each dot around the element symbolizes one of a possible 8 valence electrons. You must fill all four sides of the Element before beginning to pair electrons.  Atomic Size: The size of an atom increases going down a group, and decreases from left to right across a period  Ionization Energy: energy required to remove one valence electron. As distance from the nucleus to valence electron increases, the ionization energy decreases. In other words, Ionization energy increases going up a group and from left to right on a period, while it decreases down a group and from right to left across a period.  The Octet Rule: elements are most stable when they have 8 valence electrons in their outer shell. This is a reason that the noble gasses are so unreactive- they have 8 valence electrons. For main group elements, (1A-8A) the group number is also the number of valence electrons in that element’s outer shell.  Ionic compounds o Formed when an atom loses or gains an electron. o Cations are positive ions which lose an electron which form from metals in group 1A-3A  Naming: Element name+ “ion,” for example, calcium ion o Anions are negative ions which gain an electron which form between nonmetals in group 5A-7A  Naming: Drop ending of element name, and add "- ide,” for example, sulfide or oxide o Transition elements always have a positive charge, and can form more than one charge, except for Cd, Zn, and Ag, which have fixed charges of 2+, 2+, and 1+ respectively. A roman numeral equal to the ion’s charge is placed in parenthesis directly after the element name. For example Iron (III). o When writing ionic compounds, cross the charges of the two elements to get a correct formula  Polyatomic Ions: −¿ ¿ o nitrate:N O3 o 3−¿ ¿ phosphate:PO 4 o 2−¿ ¿ sulfate:SO4 o 2−¿ ¿ carbonate:CO 3 −¿ o bicarbonate:HCO ¿ 3 −¿ ¿ o nitrite:NO2 o 2−¿ ¿ sulfite:S3 o 3−¿ ¿ phosphite:PO 3  Notice that in ions ending in –ite contain one less oxygen than ions ending in –ate  Exceptions −¿ o hydroxide:OH ¿ +¿ o ammonium:NH ¿ 4 −¿ ¿ o cyanide:CN −¿ o acetate:CH COO ¿ 3  Each polyatomic ion formed by halogens have a -1 charge and follow the same pattern as follows: −¿ ¿ o Perchlorate: ClO 4 −¿ o Chlorate: ¿ going down the list, ClO3 −¿ o Chlorite: ClO 2 each compound has one less oxygen −¿ ¿ o Hypochlorite: ClO −¿ o Chloride: Cl ¿  Covalent Compounds: o Formed when a valence electron is shared to achieve stability between non metal compounds o The names of covalent compounds need prefixes since several compounds can be formed with the same non metals  1: mono  2: di Naming:  3: tri  4: tetra First word: only use prefix if element has a  5: penta subscript greater than one  6: hexa Second word: always gets prefix when naming  7: hepta  8: octa  9: nona  10: deca  ex: nitrogen dioxide= NO 2  ex: dinitrogen tetroxide= N2O 4  Lewis Dot Structures of Compounds: Step by step o Arrange atoms o Calculate the number of valence electrons o Place a bonding pair of electrons between each of the elements o Place remaining electrons as lone pair to complete the octets  Double and Triple Bonding: o These bonds will most likely include one of the following elements: carbon, oxygen, nitrogen, or sulfur o These bonds form when single covalent bonds fail to complete octets in all atoms  When more than one way is possible to write a Lewis Structure, the molecule is said to have resonance structure  Certain elements can only exist as a diatomic molecule, meaning that the element is always paired with another atom of that element. These 7 elements are hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine.  VSEPR Theory: Valence Shell Electron-Pair Repulsion o Describes how electrons are arranged around the central atom, and that electron groups are arranged as far apart as possible o The shape of a molecule is determined by the number of atoms attatched to the central atom.  2 electron groups: linear shape  3 electron groups: trigonal planar or bent  4 electron groups: tetrahedral, bent, or trigonal pyramidal  Nonpolar molecules contain non-polar bonds  Polar molecules occur because one end of the molecule is more negatively charged. In a covalent polar bond, electrons are shared unequally. o Attractive forces called dipole-dipole attractions help form the shape of polar molecules- think opposites attract o Hydrogen bonds are the strongest type of dipole-dipole bond because a hydrogen atom attatches to a highly electronegative atom such as fluorine, oxygen, or hydrogen  Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons CH 4-6 Practice Problems 1. What are the atomic numbers of oxygen, copper, and calcium 2. How many neutrons are in 5Co ? 27 3. How many protons and electrons are in carbon? 4. If an element has a mass number of 4 and an atomic number of 2, how many protons does it have? What element is this? 5. Silver has 2 naturally occurring isotopes. One with an abundance of 51.82% and mass of 106.9041. The other has a 48.18% abundance and a mass of 108.9047. What is the average atomic mass of these two isotopes? 6. The atomic mass of lithium is 6.941. The two naturally occurring isotopes have masses of 6.10512 ( 6L )and 7.01600 ( 7L ). ❑ ❑ Find the percent abundance of these two isotopes. 7. What is the electron configuration for sodium? 8. What is the Electron configuration for cobalt? 9. Which element has the largest atomic size in period 4? 10. Which element has the lowest ionization energy in group 7A? 11. Draw the Lewis structure for Oxygen. 12. Write the formula for aluminum oxide. 13. Write the formula for Copper (II) Nitrate. 14. Write the formula for Tin (IV) Hydroxide. 15. Write the Formula for Potassium Pentachloride. 16. Name the compound H2O . 17. Draw CH 4 in its Lewis structure. 18. Draw H 2O in its Lewis Structure. 19. Draw the Resonance structures for S O 3 Solutions 1. Oxygen: 8 Copper: 29 Calcium: 20 2. 32 3. 6 protons and 6 electrons 4. 2 protons, He 5. 107.87 amu 6. ( ❑L ) 8.2% 7 ( ❑L ) 91.8% 7. 1s 2s 2 p 3s 1 2 7 8. [Ar] 4s 3d 9. Potassium 10. Astatine 11. 12. Al2O 3 13. NO 3 3 Al¿ 14. OH¿ 4 Sn¿ 15. PCl 5 16. dihydrogen monoxide 17. 18. 19.


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