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by: Layan Hamidi Nia

EXAM 1 STUDY GUIDE Chem 10060-001

Layan Hamidi Nia

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About this Document

This study guide covers almost 95% of what is going to be on the test.. "Chapters 2, 3 & 4"
general chemistry I
margaret leslie
Study Guide
General Chemistry 1, ch 2, Ch3, Ch4, exam 1 study guide
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This 44 page Study Guide was uploaded by Layan Hamidi Nia on Thursday February 4, 2016. The Study Guide belongs to Chem 10060-001 at Kent State University taught by margaret leslie in Spring 2016. Since its upload, it has received 101 views. For similar materials see general chemistry I in Chemistry at Kent State University.


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Date Created: 02/04/16
EXAM 1 STUDY GUIDE Dalton’s Atomic Theory •All matter consists of indivisible atoms • Atoms of one kind of element are identical in mass and properties; atoms of different kinds of elements are different •Compounds are made up of definite numbers of atoms of the component elements •The weight of a compound equals the sum of the weights of the component elements Rutherford’s conducted further experiments which contradicted Thompson’s model. To explain his results Rutherford postulated Most of the mass of the atom and all its positive charge was located in a concentrated core, called the nucleus. Most of the total volume of the atom is empty in which electrons move around the positive core. Model of the Atom Since the times of Rutherford, many more subatomic particles have been discovered. However, for chemists three sub-atomic particles are all that we need to focus on – ELECTRON, PROTON, NEUTRON. Electrons are –1, protons +1 and neutrons are neutral. Atoms have an equal number of electrons and protons they are electrically neutral. Protons and neutrons make up the heavy, positive core, the NUCLEUS which occupies a small volume of the atom. Atoms of the SAME element can have different number of NEUTRONS. These atoms of the SAME elements but with different number of neutrons are called ISOTOPES. Hence isotopes of the same elements have the same number of protons and electrons, but different number of neutrons and hence different masses. EXAMPLE – Carbon has three isotopes C12, C13, C14. Each of these isotopes differ by the number of neutrons – ALL have SIX protons. C12 has SIX neutron, C13 has SEVEN and C14 has EIGHT. To denote the number of protons and neutrons in an atom the following symbol notation is used 126 where 12 denotes SUM OF PROTONS + NEUTRONS 6 denotes the number of PROTONS So for the isotopes of carbon the complete chemical symbols are: 126C, 16C, 146 The superscript, which is the sum of the number of protons and neutrons, is called the MASS NUMBER (A). The subscript indicates the number of protons and is called the ATOMIC NUMBER (Z) . How many protons, neutrons, and electrons are there in 197 79 Au atomic number and refers to the number of protons. Hence this atom has 79 electrons ande 197-79 = 118 neutrons Average Relative Atomic Mass Because the abundance of the isotopes of different elements are essentially constant, we can define an AVERAGE RELATIVE ATOMIC mass Average Relative Atomic Mass = average mass of atoms of an element = (Abundance) AMass) A (Abundance) (MBss) + A Problem 35 37 Naturally occurring chlorine has two isotopes, 17l, 17Cl. The 35-Cl isotope has a relative atomic mass of 34.9688 and an abundance of 75.77% and the 37-Cl isotope has a relative atomic mass of 36.9659 and an abundance of 24.23%. Calculate the average atomic mass of Cl. Average Atomic Mass of Cl = (0.7577x34.9659) + (.2423x36.9659) = 35.4527 Average relative atomic mass of C is 12.0107 accounting for C (98.892%, relative atomic mass 12.000000) and C (1.108%, relative atomic mass 13.003354) PROBLEM How many moles of Fe are there in 8.232 g of Fe? How many atoms are there in 8.232 g of Fe? Moles of Fe = 8.232 g Fe x 1 mole = 0.1474 mol Fe 55.85 g Fe How many grams of water are there in 0.2000 moles of water? 0.2000 mol H O x 18.015 g H O = 3.603 g H O 2 2 2 1 mol H O 2 What limits the “boundary” of a molecule? Atoms in molecules are held together by strong interactions called CHEMICAL BONDS Interactions between neutral atoms in a molecule (e.g. H , 2 H 2) is called a COVALENT bond, forming covalent compounds Interactions between charged elements (IONS) result in a different kind of chemical bond, called an IONIC BOND Compounds formed via interactions between ionic (charged) elements are called ionic compounds. ELEMENTAL ANALYSIS Determination of the relative amounts of the elements in a compound Determines the EMPIRICAL FORMULA which is the simplest possible formula and indicates the relative amounts of constituent elements For example, the molecular formula of hydrogen peroxide is H 2 2 its empirical formula is HO. The empirical and the chemical formula can be the same, for example, H 2. For ionic compounds, the empirical formula is the same as the chemical formula (NaCl) Problem An oxide of nitrogen is analyzed and found to contain 25.9% N and 74.1% O. What is the empirical formula of the compound? In 100.0 g of compound: # moles of N = 25.9 g N / (14.01 g/mol) = 1.85 mol # moles of O = 74.1 g O / (16.00 g/mol) = 4.63 mol Ratio of O: N :: 2.5:1 Must be whole numbers – N O 2 5 Determining Chemical Formulas The empirical formula tells you the simplest ratio of the individual elements in the compound. For an ionic compound this information is enough. For a molecular compound this may not be enough since the empirical formula may not be the molecular formula. Knowledge of the MOLAR MASS of the compound and its empirical formula, allows the molecular formula to be determined. Elemental analysis of a sugar shows that it consists of 40.0% carbon (C), 6.7% hydrogen (H), 53.3% oxygen (O). The molar mass of the compound was found to be 180.0 g/mol. What is the molecular formula of the compound? moles of C in 100.0 g of compound = (mass of C g)/(atomic mass g/mol) = 40.0/12.01 = 3.33 mol C in 100.0 g of compound moles of H in 100.0 g of compound = (mass of H g)/(atomic mass g/mol) = 6.7/1.01 = 6.7 mol H in 100.0 g of compound moles of O in 100.0 g of compound = (mass of O g)/(atomic mass g/mol) = 53.3/16.00 = 3.33 mol O in 100.0 g of compound Ratio of C:H:O :: 1:2:1; hence empiricalformula is CH 2 Molar mass of empirical formula = 12.01 + 2(1.01) + 16.00 = 30.0 g/mol Ratio of molar mass of compound : molar mass of empirical formula 180.0/30.0 = 6.0 molecular formula is (CH O) or C H O 2 6 6 12 6 An important aspect of a chemical reaction is that MASS IS ALWAYS CONSERVED - i.e. the total mass of the reactants must equal the total mass of the products. To ensure that mass is conserved, we have to keep track of the number of atoms of each element in the reactants and number of atoms of each element in the products Writing Chemical Equations H 2 O 2H O 2 reactants products Mass has not been conserved - equation is not BALANCED Need to make sure that the number of atoms of a given element is the same on either side. H 2 O  H 2 2 H is balanced, but O is not. To balance O, multiply H2O by 2 H2+ O 2 2 H O 2 Now O is balanced but not H (4 H’s on right, 2 on left) Multiply H by 2 2 2H 2 O 22 H O 2 Chemical Stoichiometry Reactants are consumed and products are formed in definite proportions. These proportions are given by the coefficients in the balanced equations for chemical reactions The calculation of the quantities of reactants and products is called STOICHIOMETRY. Stoichiometry is the use of chemical equations to calculate quantities of substances that take part in chemical reactions. To do a stoichiometric calculation, the chemical equation for the calculation must be balanced. Equations are read in terms of moles of reactants and product 2H (g) + O (g)  2H O(g) 2 2 2 2 moles of H (g) reacts with 1 mole of O (g) to form 2 2 2 moles of H O2g) Or: 2N oolecules of H (g) 2 N molecoles of O (g)  2N 2 o molecules of H O(g) 2 Problem Consider the reaction of 100 g of H (g) with sufficient O (g) 2 2 to produce the stoichiometric quantity of H O(g). 2 (stoichiometric quantities are the exact amounts of reactants and products predicted by balanced equations). Calculate the mass of H O formed2 Need a balanced equation for this reaction 2H 2g) + O (g2  2H O(g) 2 To find the mass of water formed, need to find the number of m2les of H that reacted: (mass of 2 ) = (100.g 2 )/(2.02 g/mol) = 49.5 m2l H (molar mass of 2 ) 2 moles of H reacts with 1 mole of O to form 2 moles of H O 2 2 2 => 49.5 moles of H will form 49.5 moles of H O 2 2 Hence, the mass of H O(g) formed 2 =(49.5 moles H O) x (18.02 g/mol) = 892 g H O 2 2 Problem: Calcium carbonate CaCO (s) is 3ecomposed by HCl(aq) to give CaCl 2 (aq), CO (2) and H O(2). If 10.0g of CaCO are tr3ated with 10.0 g of HCl, how many grams of CO are gen2rated? First write a balanced equation for the reaction: CaCO (s3 + 2HCl(aq)  CaCl (aq) + 2 O(l) + C2 (g) 2 1 mole of CaCO reacts with 2 moles of HCl to form 1 mole of CO 3 2 Moles of CaCO = (13.0 g)/(100.0g/mol) = 0.100 mole CaCO 3 Moles of HCl = (10.0 g)/(36.5 g/mol) = 0.274 mol HCl Hence 0.1 mole of CaCO react3 with 0.2 moles of HCl. Since the amount of HCl present is greater that 0.2, HCl is in excess and CaCO is the limiting reactant. 3 Hence, 0.1 mole of CO form2d. Mass of CO formed = 0.1 mol x 44.01 g/mol = 4.40 g CO 2 2 Product Yields In the previous calculation 4.40 g is the amount of CO we would expect to be formed 2 4.40 g CO 2s the CALCULATED or the THEORETICAL PRODUCT YIELD. This assumes that the reaction goes to completion, and that there are no competing factors that may reduce the amount of CO formed. 2 The measured amount of product formed is called the ACTUAL YIELD which is often smaller than the theoretical yield. Percentage Yield = (actual yield) x 100 % (theoretical yield) The larger the % yield the more cost effective is the process and hence a more likely candidate for industrial scale processes (other factors are also important: the nature of the by-products - are they environmentally safe -, the cost of the starting materials, etc.) Chemical Bonding A chemical bond results from strong electrostatic interactions between two atoms. The nature of the atoms determines the kind of bond. COVALENT bonds result from a strong interaction between NEUTRAL atoms Each atom donates an electron resulting in a pair of electrons that are SHARED between the two atoms For example, consider a hydrogen molecule, H . W2en the two hydrogen, H, atoms are far apart from each other they do not feel any interaction. As they come closer each “feels” the presence of the other. The electron on each H atom occupies a volume that covers both H atoms and a COVALENT bond is formed. Once the bond has been formed, the two electrons are shared by BOTH H atoms. An electron density plot for the H m2lecule shows that the shared electrons occupy a volume equally distributed over BOTH H atoms. Electron Density for the H2molecule What factors determine if an atom forms a covalent or ionic bond with another atom? The number of electrons in an atom, particularly the number of the electrons furthest away from the nucleus determines the atom’s reactivity and hence its tendency to form covalent or ionic bonds. These outermost electrons are the one’s that are more likely to “feel” the presence of other atoms and hence the one’s involved in bonding i.e. in reactions. Chemistry of an element depends almost entirely on the number of electrons, and hence its atomic number. Columns are called GROUPS (FAMILIES) and rows are called PERIODS. Elements in a group have similar chemical and physical properties. The total number of electrons within a group is different, increasing in number down a group However, the number of electrons furthest away from the nucleus, called the OUTER or VALENCE electrons is the same for all elements in a group. Groups are referred to by names, which often derive from their properties I – Alkali metals; II – Alkaline Earth metals VII – Halogens; VIII – Noble gases The elements in the middle block are called TRANSITION ELEMENTS ELECTRONEGATIVITY measures the tendency of one atom to attract electrons from another atom to which it is bonded. For example, Metallic elements loose electrons (to form positive ions) more readily than non-metallic elements Metallic elements are hence referred to as being more ELECTROPOSITIVE that non-metals. Non-metals are more ELECTRONEGATIVE compared to metals Based on the position of elements in the periodic table, we can determine the kind of bond formed Generally: Nonmetallic element + nonmetallic element  Molecular compound Molecular compounds are typically gases, liquids, or low melting point solids and are characteristically poor conductors. Examples are H O, 2H , NH4. 3 Generally, Metallic compound + nonmetallic compound  IONIC compound Ionic compounds are generally high-melting solids that are good conductors of heat and electricity in the molten state. Examples are NaCl, common salt, and NaF, sodium fluoride . NAMING COMPOUNDS The chemical formula represents the composition of each molecule. In writing the chemical formula, in almost all cases the element farthest to the left of the periodic table is written first. So for example the chemical formula of a compound that contains one sulfur atom and six fluorine atoms is SF . 6 If the two elements are in the same period, the symbol of the element of that is lower in the group (i.e. heavier) is written first e.g. 3F . In naming covalent compounds, the name of the first element in the formula is unchanged. The suffix “-ide” is added to the second element. Often a prefix to the name of the second element indicates the number of the element in the compound SF 6 sulfur hexafluoride P O – tetraphosphorous decoxide 4 10 CO – carbon monoxide CO 2 carbon dioxide Electrolytes & nonelectrolytes battery bulb + Na + + - Cl- deionizedwater+ NaCl NaCl(s) --> Na (aq) + Cl (aq) + When the battery is turned on the Na ions flow toward the negative plate (anode) and the Cl ions to the positive plate (cathode). The flow of ions constitutes a current. The circuit is now complete, current flows through the circuit, and the bulb turns on. NaCl is called an electrolyte. Electrolyte : a compound which when dissolved in a solvent dissociates to form ions in solution. Typically electrolytes are ionic compounds since they dissolve in solution to form ions. Example: K S2 4 Some covalent compounds (like acids and bases) can dissociate in solution to form ions. Electrolytes are characterized as being strong or weak. The strength of an electrolyte depends on the degree to which the compound dissociates in water to form ions. Hence ionic compounds like NaCl and K SO wh2ch 4 dissociate completely in water are strong electrolytes. Weak electrolytes do not dissociate extensively in water- consequently the conductance of a solution of a weak electrolyte in low. Non-electrolytes do not dissociate in solution to form ions and hence their solutions do not conduct electricity.


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