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CHEM 101 - General Chemistry 1 Chapter 2 – Atoms and Elements Handout prepared by Dr.Saravanan Rajendrasozhan, Department of Chemistry, University of Hail. Assistant Professor, Reference book: ‘Principles of Chemistry: A Molecular Approach, First edition, by Nivaldo Tro. Published by Pearson Education International. MATTER Matter is anything that has both mass and volume (occupies space). It can be a solid, liquid or gas. Early Philosophy of Matter ¾ Because of the lack of sophisticated instruments in the early days, there are two different opinions given by philosophers about matter. 1) Matter had an ultimate, tiny, indivisible particle (Leucippus and Democritus) 2) Matter was infinitely divisible (Plato and Aristotle) ¾ In the late 17th century, research on matter was improved that could not explain the infinitely divisible matter concept. Atoms ¾ Atoms are basic unit of matter or building blocks of matter. Example: Carbon (C), Oxygen (O), Hydrogen (H), Gold (Au) Elements ¾ A pure chemical substance consisting of one type of atom. Example: Pure gold bar (made of only gold atoms) ¾ Atoms are the smallest particle of the element which retains the properties of that element. Pure gold bar (element) Gold bar (element) consisting of only gold atoms Gold atom (smallest particle of the element) FUNDAMENTAL CHEMICAL LAWS Scientific law is a description of an observed phenomenon. Law of Conversion of Matter ¾ In a chemical reaction, mass (weight) is neither created nor destroyed. Total mass of reactants = Total mass of products Reactants Products Reaction 2Na (s) + Cl2 (g) 2NaCl (s) Mass of chemicals 7.7 g Na + 11.9 g Cl2 19.6 g NaCl Total mass 19.6 g 19.6 g Law of Definite Proportions A given compound always contains exactly the same proportion of elements by mass (regardless of their source or how they were prepared). Sample 1 100 g Sodium chloride salt (NaCl) produced by reacting sodium metal and chlorine gas Sample 2 58.44 g of NaCl prepared by seawater evaporation process Elements: 39.3 g Sodium and 60.7 g Elements: 22.99 g Sodium and chlorine 35.44 g chlorine Proportion=Mass of Cl/Mass of Na Proportion=Mass of Cl/Mass of Na = 60.7 g / 39.3 g = 35.44 g / 22.99 g = 1.54 = 1.541 Law of Multiple Proportions ¾ When two elements (A and B) combine to form two different compounds, they combine in a ratio of small whole numbers. ¾ The masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers. Example: Carbon and Oxygen can combine at different ratio to form carbon monoxide (CO) and carbon dioxide (CO2). CHEM 101‐Chapter 2 Handout Page 2 CO: 1 g carbon combine with 1.33 g oxygen CO: 0.752 g carbon combine with 1 g oxygen CO2: 1 g carbon combine with 2.67 g oxygen CO2: 0.375 g carbon combine with 1 g oxygen Ratio of oxygen combine with 1 g carbon = 2.67 g/1.33 g Ratio of carbon combine with 1 g oxygen = 0.752 g/0.375 g =2 =2 Dalton’s Atomic Theory Scientific theory is an explanation of the observed phenomenon. 1) Each element is composed of tiny, indestructible particles called atoms. 2) All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements. 3) Atoms combine in simple, whole-number ratios to form molecules of compounds. 4) In a chemical reaction, atoms of one element cannot change into atoms of another element. During a reaction, the atoms are simply rearranging the way they are attached. EARLY EXPERIMENTS TO CHARACTERIZE ATOMS ¾ Three experiments are considered important in characterizing atoms. 1) Thomson’s Experiment – Postulated the existence of electrons. 2) Millikan’s Oil Drop Experiment – Measured the mass and charge of electrons. 3) Rutherford’s Experiment – Proposed the nuclear model of atom. Thomson’s Experiment ¾ Glass tube (almost all the air has been evacuated) was connected to metal electrodes. ¾ When high voltage is applied, the cathode starts to emit a glowing ray, which is called cathode ray. ¾ When electric field applied around the glass tube, the cathode ray deflects towards the positively charged plates. ¾ The ray also deflects when magnetic field is applied. - From the observation, Thomson postulated that cathode ray contains some particles with negative charge and mass. CHEM 101‐Chapter 2 Handout Page 3 - Also calculated the charge/mass ratio of the particle: −1.76 × 108 C/g ¾ These cathode ray particles became known as electrons. Thomson’s Plum Pudding Model of Atom ¾ The structure of the atom contains many negatively charged small particles, called electrons. ¾ The electrons are held within a positively charged sphere (positively charged cloud). ¾ This model of atom is called as palm-pudding model (Later experimetns proved that this model is not correct). Millikan’s Oil Drop Experiment ¾ The apparatus contains 2 horizontal metal plate electrodes (anode and cathode). ¾ A fine mist of oil droplets was sprayed (by atomizer) into a chamber above the plates. ¾ When the oil droplets enter the space between the electrodes (through the hole in plate electrode), the oil droplets were ionized by applying ionizing radiation such as x-ray. ¾ The ionized droplets could be made to rise and fall by changing the voltage across the electrodes (anode and cathode). ¾ From this experiment, the mass and charge of electron was calculated. Mass of electron = 9.1 X 10-28 g Charge of electron = −1.60 × 1019 C. Radioactivity ¾ The discovery radioactivity initiated further experiments to characterize the atomic structure. ¾ Certain elements would constantly emit small, energetic particles and rays. This is called radioactivity. ¾ These particles and rays can penetrate matter. Rays Charge Form α: Alpha particles Positive charge Particles with a mass of 4H atoms β: Beta particles Negative charge Particles with a mass ~1/2000 H atom γ: Gamma Rays No charge Energy rays (not particles) Rutherford’s Experiment ¾ Alpha particles (positively charged) were directed at an ultra-thin sheet of gold foil. ¾ The deflection in alpha rays was monitored. # Results Conclusion 1 98% of α-particles penetrate gold foil and went straight through. An atom is mostly empty space. 2 0.005% of α-particles bounced back from the gold foil. An atom contains a dense particle (small in volume compared to the atom but large in mass). 2 2% of the α-particles penetrate gold foil, but deflected by large angles. The dense particle in the atom is positively charged which deflect positively charged α-particles Rutherford’s Nuclear Model of Atom ¾ The atom has a tiny dense center called nucleus. It is responsible for the entire mass of an atom and it has positively charged particles (named protons). ¾ Most of the volume of atom is empty space. The negatively charged electrons are dispersed in the empty space surrounding the nucleus. ¾ The atoms are neutral as it has equal number of positively charged protons and negatively charged electrons. ¾ Later, Rutherford and his student J. Chadwick discovered the presence of neutrons (no charge) in the nucleus of the atom. Protons and neutrons are collectively responsible for the mass of an atom. STRUCTURE OF ATOM ¾ Thus, an atom consist of 1) Protons 2) Neutrons 3) Electrons ¾ Protons and neutrons are present in the tiny nucleus at high density. Both of them accounts for almost all the atom’s mass (atomic mass). ¾ Electrons form the outer shell of the atom, so that it allows one atom to interact with another. Thus, it is responsible for the chemical properties of the atom. ¾ Atoms are neutral in charge. So they must have equal number of positively charged protons and negatively charged electrons. Subatomic Particle (& Charge) Proton Location Charge +1 1.6726 x10-24 Determine the identity of an atom. Nucleus 0 1.6749 x10-24 Prevent the repulsion between positively charged protons and help to pack the protons in nucleus. Move around nucleus -1 0.0009 x10-24 Determine the size and chemical properties of the atom. (No charge) Electron (Negative) Role Nucleus (Positive) Neutron Mass (g) CHEM 101‐Chapter 2 Handout Page 6 Changes in the Structure of Atom ¾ The number of protons and neutrons in an atom cannot be changed by the chemical reaction. ¾ The number of electrons in an atom may be changed during the chemical reaction to form a charged atoms (positively or negatively charged ions). ELEMENTS and Number of Protons ¾ The elements were identified by the number of protons in the nucleus. ¾ To identify the different elements (atoms), we have a standard way of representing each element: Chemical symbol with atomic number and mass number. Chemical Symbol: One or two letter abbreviation (One capital letter or one capital letter with one lower case letter). It is based on current name or old Latin name. Current Name Latin Name Symbol Hydrogen - H Carbon - C Calcium - Ca Oxygen - O Iron Ferrum Fe Copper Cuprum Cu Silver Argentum Ag Sodium Natrium Na Atomic number (Z) = Number of protons Mass number (A) = Number of protons + Number of neutrons ¾ Atomic number (Z) is written as a subscript to the left of chemical symbol. ¾ Mass number (A) is written as a superscript on the left of chemical symbol. ¾ Z is always smaller than A (except for Hydrogen). M A Example: N N Carbon Chemical symbol = C, Atomic number = 6, Mass number = 12 The element (carbon) is written as: CHEM 101‐Chapter 2 Handout Page 7 ¾ Surprisingly in the periodic table, atomic number (Z) is mentioned on top of the chemical symbol. Example: Carbon is written as 6 C Exercise: How do you symbolize an element with the atomic number of 12 and mass number of 24. Z = 12, A = 24, From the periodic table: Element with the atomic number 12 = Mg Number of Proton, Neutron and Electrons Number of protons = Atomic number (Z) Number of electrons = Number of protons (only for neutral atoms) Number of neutrons = Mass number - Number of protons (because, Mass number = Number of protons + Number of neutrons) Example: How many protons, electrons, and neutrons are in an atom of ? Atomic number (Z) = 24 Mass number (A) = 52 Number of protons = 24 Number of electrons = 24 Number of neutrons = 52 – 24 = 28 Exercise: Calculate the number of proton, electron and neutrons in the following atoms: Atom # of Protons # of Electrons # of Neutrons Pb 82 82 208-82 = 126 F 9 9 19-9 = 10 Ca 20 20 41-20 = 21 U 92 92 235-92 = 143 ISOTOPES and Number of Neutrons ¾ Atoms with the same atomic number but different mass numbers are called isotopes. ¾ Isotopes have same number of proton and electron, but different number of neutron. CHEM 101‐Chapter 2 Handout Page 8 ¾ As the chemical property of the atoms depends mainly on electrons, the isotopes have same chemical property. Example: Sodium ( Na) has 11 electrons and 11 protons. If sodium has 12 neutrons, A = 11+12 = 23, it is written as Na or Na-23. If sodium has 13 neutrons, A = 11+13 = 24, it is written as Na or Na-24. 9 Na (Na-23) and Na (Na-24) are called as isotopes. CHARGED ATOMS and Number of Electrons ¾ Atoms or group of atoms (molecules) with a net positive or negative charge are called ions. Example: Na+, NH4+, Cl–, OH– ¾ Ions are classified into 2 groups based on the number of atoms. 1) Mono-atomic ions (Example: Na+, Cl–) 2) Poly-atomic ions (Example: NH4+, OH–) The net charge of each atom is equal to the charge of poly-atomic ion. ¾ Ions are classified into 2 groups based on the charge. 1) Positive ions, called cations (Example: Na+, NH4+) 2) Negative ions, called anions (Example: Cl–, OH–) ¾ Only the number of electrons is changed in ions as compared to the parent atoms (NO change in the number of protons or neutrons). Cations ¾ When atoms or molecules lose one or more electrons, they become positively charged ions, called cations. ¾ The name of the cation is same as the element. Example 1: Sodium ion ( Na+) formation from sodium ( Na) atom. Na Lose one electron ⇓ 11 Protons, 11 Electrons Net charge = +11-11 = 0 Neutral Na+ ⇓ 11 Protons, 10 Electrons Net charge = +11-10 = +1 Positive charge Example 2: Calcium ion ( Ca2+) formation from calcium ( Ca) atom. Ca ⇓ 20 Protons, 20 Electrons Lose two electrons Ca2+ ⇓ 20 Protons, 18 Electrons Anions ¾ When atoms or molecules gain one or more electrons, they become negatively charged ions, called anions. ¾ Anions are named by changing the ending of the element name to -ide. Example 1: Chloride ion ( Cl–) formation from chlorine atom ( Cl). Cl Gain of one electron ⇓ 17 Protons, 17 Electrons Net charge = +17-17 = 0 Neutral Cl– ⇓ 17 Protons, 18 Electrons Net charge = +11-18 = -1 Negative charge Example 2: Oxide ion ( O2–) formation from oxygen ( O) atom. O O2– Gain two electrons ⇓ 8 Protons, 8 Electrons ⇓ 8 Protons, 10 Electrons Name of common monoatomic anions: Name of Atom Name of Anion Name of Atom Name of Anion H: Hydrogen H– : Hydride O: Oxygen O2– : oxide F: Fluorine F– : Fluoride S: Sulfur S2– : Sulfide Br: Bromine Br– : Bromide N: Nitrogen N3– : Nitride I: Iodine I– : Iodide P: Phosphorus P3– : Phosphide Cl: Chlorine Cl– : Chloride Number of Proton, Neutron and Electrons in Ions ¾ The number of protons and neutrons in an ion is same as that of the parent atom. Example: The number of protons and neutrons in sodium ion (Na+) is same as that of sodium atom (Na). ¾ The number of electrons in an ion may be increased or decreased as compared to number of proton (atomic number). Number of electrons in an ion = Atomic number – Charge Exercise: Calculate the number of proton and electron in the following ions: Ions Z (from Periodic table) # of Protons # of Electrons Zn 2+ 30 30 30-(+2) = 28 N3– 7 7 7-(-3) = 10 Br– 35 35 35-(-1) = 36 Ions and Compounds ¾ Ions behave much differently than the neutral atom. Example: Sodium metal (Na) is highly reactive and quite unstable. Sodium ion (Na+) found in table salt (NaCl) is non-reactive and stable. ¾ Cations (positive charged ions) and anions (negative charged ions) attract each other because of the opposite electrical charges. ¾ Cation and anion with equal amount of positive and negative charges join together to form a compound, because compounds are neutral (Net charge=0). 9 Cation: Na+ (charge: +1) 9 Anion: Cl– (charge: -1) 9 Compound: 1Na+ + 1Cl– → NaCl ¾ If the cation and anion has different amount of positive and negative charges, multiple ions join together to make equal positive and negative charges to form a compound. 9 Cation: Mg2+ (charge: +2), Anion: Cl– (charge: -1), Compound: MgCl2 9 Cation: Fe3+ (charge: +3), Anion: O2– (charge: -2), Compound: Fe2O3 CHEM 101‐Chapter 2 Handout Page 11 PERIODIC TABLE ¾ All the known elements are arranged in periodic table depending on the atomic number (Z) and chemical properties. ¾ Each row (left to right) of the periodic table is called ‘Period’. ¾ Each column (top to bottom) of the periodic table is called ‘Group’. 9 Atomic number is mentioned on top of the chemical symbol of element. It is increased from left to right of periodic table (period). 9 The chemical properties of elements are similar in each group (top to bottom of periodic table). ¾ From the periodic table you can find out number of protons and number of electrons present in an atom. Example: 2 Atomic number (Z) = 2 He Number of protons = 2 Number of electrons = 2 Periodic Pattern ¾ Periodic pattern are specific trends that are present in the periodic table, which help to predict the properties of elements. ¾ The main periodic trends include: metallic character, electronegativity, ionization energy, electron affinity and atomic radius. CHEM 101‐Chapter 2 Handout Page 12 Pattern in Metallic Character ¾ Metallic character (ability to lose electrons) decreases from left to right of the periodic table. ¾ Metallic character increases from top to bottom of the periodic table. ¾ Based on the metallic properties, the elements are classified into: 1) Metals, 2) Non-metals, and 3)Metalloids Metals ¾ Metals conduct heat and electricity. ¾ Metals are solids at room temperature (except Hg). They are malleable (can be shaped) and ductile (pulled into wires). ¾ Most of the elements (about 75%) are metals. ¾ Metals frequently lose electrons to become positive ions (cations). ¾ It can form ionic compound with a non-metal ion (negatively charged anion). ¾ Metals are in the left side and middle of the periodic table. Non-metals ¾ Nonmetals are poor conductors of heat and electricity. ¾ Nonmetals found in all three states (solid, liquid, gas) at room temperature. Solids are brittle (easily breakable), so they are NOT malleable and NOT ductile. CHEM 101‐Chapter 2 Handout Page 13 ¾ Non-metals frequently gain electrons to become negative ions (anions). ¾ They can form ionic compound with a metal ion (positively charged cation). ¾ They can also form covalent compound by bonding with another nonmetal. ¾ Non-metals are in the right side of the periodic table (diagonal from B to At) and H. Metalloids ¾ Metalloids show some properties of metals and some properties of nonmetals. ¾ They are semiconductors of heat and electricity. ¾ They are present in between the metals and nonmetals in periodic table. MODERN PERIODIC TABLE ¾ Elements with similar chemical and physical properties are in the same column. Remember: Columns are called Groups. Rows are called Periods ¾ They are classified into 3 groups: 1) Main group elements 2) Transition elements 3) Inner-transition elements. 1) Main group elements CHEM 101‐Chapter 2 Handout Page 14 9 Also called as Representative elements or Group A elements (present in Group 1A to 8A). 9 They are both metals and non-metals. 2) Transition elements 9 Also called as Group B elements (present in Group 1B to 8B) 9 They are all metals. 3) Inner-transition elements 9 Also called as Rare Earth Elements (present in bottom rows, really belongs to periods 6 and 7) 9 They are all metals. Important Groups ¾ Important groups in the periodic table are: Hydrogen, Alkalai metals, Alkalai earth metals, Halogens and Noble gases. Hydrogen ¾ Nonmetal present on top of the Group 1A. ¾ Colorless, diatomic (H2) gas with very low melting point and low density. 9 Reacts with nonmetals to form molecular compounds (Example: HCl and H2O) ¾ Reacts with metals to form hydrides. Metal hydrides react with water to form H2 gas. ¾ HX dissolves in water to form acids. Alkalai Metals ¾ Metals present in group 1A (Li, Na, K, Rb, Cs, Fr) react with water to form alkaline solutions (pH>7, basic). So they are called alkaline metals. Example: 2 Na + 2 H2O 2 NaOH + H2 + Heat ¾ Hydrogen is also present in Group 1A, but it is a nonmetal (NOT belongs to alkalai metals). ¾ They are soft solids with low melting point and low density. ¾ They are very reactive. 9 React with nonmetals to form water-soluble compounds. Example: 2Na + Cl2 2 NaCl CHEM 101‐Chapter 2 Handout Page 15 Alkalai Earth Metals ¾ Metals present in group 2A (Be, Mg, Ca, Sr, Ba, Ra) react with oxygen to form metal oxides, which can dissolve in water to form alkaline solution. Example: 2Mg + O2 2MgO All these metals are found in Earth and remain solids (earths) in fire. So they are called alkaline earth metals. ¾ They are hard solids with very high melting point and high density. ¾ They are very reactive, but the reactivity is lesser than alkaline metals (Group 1A). ¾ Alkaline earth metals (except Be-beryllium) react with water to form hydrogen. Example: Ba + 2H2O Ba(OH)2 + H2 Halogens ¾ Nonmetals present in group 7A (F, Cl, Br, I, At). In Greek, ‘hal’ means ‘salt’ and ‘gen’ means ‘to produce’. Non-metals in Group 7A react with sodium to produce salt of similar properties. So they are called Halogens. Example: Cl2 + 2Na 2NaCl ¾ Exist as diatomic molecules (F2 and Cl2 are gases, Br2 is liquid, and I2 is solid). ¾ Very reactive. Reacts with metals to form ionic compounds. ¾ All HX are acids (Acidity: HI < HCl < HBr < HI) ¾ Cl2, Br2 react slowly with water. Example: Br2 + H2O HBr + HOBr Noble Gases ¾ Nonmetals present in group 8A (He, Ne, Ar, Kr, Xe, Rn) are called noble gases. ¾ They are all gases at room temperature. They have very low melting and boiling points. ¾ They are most stable or very unreactive (practically inert). Also very hard to produce ions (very hard to remove electron from or give an electron to noble gas). ¾ They are odorless and colorless, and are used in many conditions when a stable element is needed to maintain a safe and constant environment. CHEM 101‐Chapter 2 Handout Page 16 IONIC CHARGE AND PERIODIC TABLE ¾ Different ions may have different magnitude of charge. Example: Na+, Ca2+, Fe3+, Cl–, O2– ¾ Charge on an ion can be determined from an element’s position on the periodic table. ¾ Metals always form positively charged cations. 9 For the main group metals (Groups 1A, 2A and Al from Group 3A) Charge of Cation = Group number ¾ Nonmetals form negatively charged anions. 9 For many main group nonmetals (Groups 7A, 6A and N from Group 5A) Charge of Anion = Group number – 8 ¾ Charge on transition metal ion cannot be determined from its position in the periodic table. ¾ Some transition metals form only one cations, but some transition metals (and some metals form Groups 3A and 4A) form more than one cations. Example: Ag can form only Ag+, but Fe can form Fe2+ and Fe3+ Transition metal Charge of Ion Element Charge of Ion Ag +1 Cu +1 or +2 Zn, Cd +2 Hg Hg2+, Hg22+ Cr, Mn, Fe, Co +2 or +3 Sn, Pb +2 or +4 CHEM 101‐Chapter 2 Handout Page 17 ATOMIC MASS ¾ Atomic mass (atomic weight) is the mass of an atom in atomic mass units (amu). ¾ Atomic mass can be determined using an instrument called ‘Mass Spectrometer’. ¾ Atomic mass can also be derived from the mass number of the atom. 9 Mass number (A) = Number of Protons + Number of Neutrons (Remember: It is indicated in the superscript in left side of the chemical symbol or in the bottom of chemical symbol in periodic table). Atomic mass = Mass number in amu Example: Atomic mass of carbon ( C) = 12 amu Average Atomic Mass ¾ Many elements present in nature as mixtures of isotopes (same atomic number but different mass number). Example: Carbon C (C-12) 98.89% present as 1.11% present as C (C-13) < 0.01 % present as C (C-14) ¾ Average atomic mass of an element is calculated from the atomic mass of various isotopes of the elements and its natural abundance (percentage of isotope present in the nature). Average Atomic mass of C = (12 x 98.89) + (13 x 1.11) + (14 x 0.01) 100 = (1186.68 + 14.43 + 0.14)/100 = 12.0125 = 12.01 amu ¾ Average atomic mass of element is listed directly beneath the element’s symbol in periodic table. Example: Atomic Number (Z) → 6 C Atomic Mass → 12.01 ← Chemical Symbol or Element’s Symbol COUNTING ATOMS BY MOLES ¾ If we can find the mass of a particular number of atoms, we can convert the mass of atoms into number of atoms in the sample. ¾ The unit ‘mole’ is used to find out the mass of particular number of atoms. CHEM 101‐Chapter 2 Handout Page 18 Mole ¾ Mole (also written as mol) is the unit of measurement used to measure the number of things, usually atoms or molecules. 1 mole = 6.022 x 1023 thing (atom or molecule) 6.022 x 1023 is Avogadro's number. Example: 1 mole of carbon = 6.022 x 1023 carbon atoms Exercise 1: Calculate the number of atoms in 1.9 moles of gold. 1 mol gold = 6.022 x 1023 gold atoms 23 Number of gold atoms = 1.9 mol gold x 6.022 x 10 atoms 1 mol gold = 11.44 x 1023 atoms = 1.144 x 1024 atoms Exercise 2: A copper coin contains 1.7 x 1026 atoms. How many moles of copper are there in the coin? 1 mol copper = 6.022 x 1023 copper atoms 1 mol Moles of copper = 1.7 x 1026 atoms x 6.022 x 1023 atoms = 0.28 x 103 mol = 2.8 x 102 mol Molar mass ¾ The mass of one mole of atoms is called the molar mass. ¾ Molar mass of an atom = Mass number or Atomic mass in grams Example 1: What is the molar mass of Molar mass of C? C = 12 grams Thus, Mass of 1 mole (6.022x1023) of C atoms = 12 grams Example 2: What is the mass of carbon (C) atoms? Atomic mass of C (from Periodic Table) = 12.01 amu Molar mass of C = 12.01 gram Substance Chemical Mass of 1 atom Atoms in 1 mole Symbol (atomic mass)* Mass of 1 mole (molar mass) Hydrogen H 1.008 amu 6.022 x 1023 atoms 1.008 g Carbon C 12.01 amu 6.022 x 1023 atoms 12.01 g Sulfur S 32.06 amu 6.022 x 1023 atoms 32.06 g *Atomic mass is taken from the Periodic Table CHEM 101‐Chapter 2 Handout Page 19 Exercise 1: Calculate the number of moles in 7.2 grams of sodium metal. Atomic mass of sodium (Na) = 22.99 amu (Remember: Atomic mass is mentioned in the periodic table at the bottom of chemical symbol) Molar mass = 22.9 g Therefore, 1 mol sodium = 22.99 g Moles of sodium = 7.2 g sodium x 1 mol 22.9 g sodium = 0.314 moles Exercise 2: How many carbon atoms present in 0.265 g of pencil lead. Atomic mass of carbon = 12.01 amu 1 mole carbon = 12.01 g (Molar mass) 1 mole = 6.022 x 1023 atoms Number of carbon (C) atoms = 0.265 g C x 1 mol 12.01 g C x 6.022 x 1023 atoms 1 mol = 0.133 x 1023 atoms = 1.33 x 1022 atoms CHEM 101‐Chapter 2 Handout Page 20 REMEMBER Atomic number = Number of protons Mass number = Number of protons + Number of neutrons Written symbol Mass Number → Atomic Number → Periodic table Atomic Number → ← Chemical …..Symbol Atomic Mass → ← Chemical …..Symbol .003 Number of protons = Atomic number Number of electron in neutral atom = Number of protons = Atomic number Number of electrons in ion = Atomic number – Charge of the ion Number of neutron = Mass number – Atomic number Cations = Metal ions = Ions with positive charge Anions = Non-metal ions = Ions with negative charge Charge of cations (+ve) = Group number Charge of anions (-ve) = Group number – 8 Atomic mass = Mass of 1 atom = Mass number in amu (atomic mass unit) Molar mass = Mass of 1 mole of atoms = Mass number or Atomic mass in grams Avogadro number = 6.022 X 1023 1 mole = 6.022 X 1023 units (atoms or molecules) Mole = Mass (g) / Molar mass (g/mol) Mass (g) = Mole X Molar mass (g/mol) Number of atoms = Mole X 6.022 X 1023 Mole = Number of atoms/6.022 X 1023 HOMEWORK From the book ‘Principles of Chemistry: A Molecular Approach, First edition, Nivaldo Tro. Page # 33 to 37 Exercises: 24, 26, 30, 34, 36, 40, 45, 48, 50, 56