General Chemistry 1st Year Student PDF

Faculty of Education- specific programme General Chemistry st 1 year student Mathematics department First semester 2024-2025 1  Contents Chapter 1: Matter & basic definitions Chapter 2: Unit conversion- sign...

Faculty of Education- specific programme General Chemistry st 1 year student Mathematics department First semester 2024-2025 1  Contents Chapter 1: Matter & basic definitions Chapter 2: Unit conversion- significant figures- mole concept Chapter 3: Atomic structure & Atomic theories& Quantum number Chapter 4 : Periodic table  Chapter 1 Matter & basic definitions What is chemistry What is matter States of matter Properties of matters The Kinetic Molecular Theory of Matter What is Chemistry? Chemistry is the study of composition, structure, properties, reaction and behavior of matter and the changes it undergoes. 1.Organic Chemistry: This branch of chemistry studies the structure, properties, and reactions of organic compounds, which contain carbon atoms. Organic chemistry is central to the study of substances like hydrocarbons, alcohols, carbohydrates, proteins, and many other compounds found in living organisms. 2.Inorganic Chemistry: Inorganic chemistry deals with the study of inorganic compounds, which do not contain carbon-hydrogen (C-H) bonds. It includes the study of minerals, metals, and coordination compounds. 3.Physical Chemistry: Physical chemistry combines principles of physics and chemistry to understand the physical and chemical properties of matter. It explores concepts like thermodynamics, quantum mechanics, and kinetics. 4.Analytical Chemistry: Analytical chemistry focuses on the methods and techniques used to identify and quantify matter. It includes techniques like spectroscopy, chromatography, and electrochemistry to analyze the composition of substances. 5.Biochemistry: Biochemistry explores the chemical processes and substances that occur within living organisms. It delves into topics such as metabolism, enzyme kinetics, and the structure and function of biological molecules. 5. Environmental Chemistry: This branch of chemistry is concerned with the study of chemical processes that occur in the environment. It investigates issues such as pollution, the fate of chemicals in ecosystems, and environmental remediation. 6. Theoretical Chemistry: Theoretical chemists use mathematical models and computer simulations to study chemical phenomena and predict the behavior of molecules and atoms. 7. Industrial Chemistry: Industrial chemistry is the application of chemistry to large-scale manufacturing and production processes. It involves the development and optimization of chemical processes for industrial purposes. 8. Materials Chemistry: This branch focuses on the synthesis, structure, and properties of materials, such as polymers, ceramics, and nanomaterials, and their applications in various industries. 9. Nuclear Chemistry: Nuclear chemistry is concerned with the study of nuclear reactions and the behaviour of atomic nuclei. It plays a crucial role in fields like nuclear power and radiopharmaceuticals. 10. Medicinal Chemistry: Medicinal chemistry involves the design, synthesis, and development of pharmaceutical agents and drugs. Medicinal chemists work to discover and optimize compounds for medical use. 5 What is matter?  NaCl compound Salt water homogeneous mixture Iron element sugar compound air homogeneous mixture helium element water compound salad heterogeneous mixture  Element Compound Molecules containing Atoms different two Molecules atoms or more C, Al, Ti H2O, H2SO3 Polyatomic Diatomic S8, O3 O2 , H2 8 States of matter The difference between the states is the distance between the molecules. 1.Solid: Solids have a fixed shape and volume. The particles in a solid are tightly packed and vibrate in fixed positions. They have a definite geometric structure, and the intermolecular forces are strong. Solids are characterized by their rigidity and resistance to compression. 2.Liquid: Liquids have a definite volume but take the shape of their container. The particles in a liquid are still close together but have more freedom of movement than in a solid. They can flow past each other, giving liquids their ability to pour and take the shape of their container. 3.Gas: Gases have neither a definite shape nor a definite volume. The particles in a gas are widely separated and have high kinetic energy, leading to constant motion. Gases can expand to fill the entire volume of their container, and they are highly compressible. 4.Plasma: Plasma is a unique and less commonly encountered state of matter. It consists of highly energized, ionized particles, such as free electrons and positive ions. Plasmas do not have a definite shape or volume and can conduct electricity. Stars, lightning, and some types of flames are examples of naturally occurring plasmas Physical properties of matter: Can be measured and observed without changing the composition or identity of a substances Here are some common physical properties of matter: 1.Mass: Mass is a measure of the amount of matter in an object. It is a fundamental property and is usually measured in grams or kilograms. 2.Volume: Volume is the amount of space that matter occupies. It is typically measured in cubic meters, liters, or other volume units. 3.Density: Density is the mass of an object per unit volume. It indicates how tightly packed the matter is within an object. The formula for density is: Density (ρ) = Mass (m) / Volume (V). 4.State: Matter can exist in various states, including solid, liquid, gas, and plasma, depending on its temperature and pressure. Color, Odor, Taste, Texture, Melting Point, Boiling Point, Solubility, Electrical Conductivity, Thermal Conductivity, Magnetism, Hardness, Transparency, Brittleness 10 Chemical properties of matter: is a property when the matter undergoes a chemical change or reaction 1.Reactivity: Reactivity is the tendency of a substance to undergo chemical reactions with other substances. Highly reactive materials readily react with other elements or compounds. For example, sodium is highly reactive and will react violently with water. 2.Flammability: Flammability is a substance's ability to ignite and burn in the presence of oxygen. Some materials are highly flammable, while others are non-flammable or relatively fire-resistant. 3.Corrosion: Corrosion is the gradual deterioration of a material due to its reaction with its environment. For example, iron rusts when it reacts with oxygen and water. 4.Oxidation: Oxidation refers to the gain of oxygen by a substance or the loss of electrons during a chemical reaction. Rusting is an example of oxidation, as iron combines with oxygen to form iron oxide. Acidity/Basicity (pH),Toxicity, Stability, Redox Potential, Hydrolysis, Combustion, Polymerization, Sensitivity to Light 11 The Kinetic Molecular Theory of Matter The postulates of the KMT: 1)Matter is composed of tiny particles (atoms, molecules or ions). 2)These particles are in constant, random motion and possess kinetic energy. 3)The particles also have potential energy due to intermolecular attractions. 4)The average kinetic energy increases as the temperature increases. 5)Energy is transferred from one particle to another during collisions, but at constant temperature the total energy of the system is constant. Kinetic energy is energy associated with motion; KE = ½mv2. If two particles of different masses are moving at the same speed, the heavier particle will have higher kinetic energy. Similarly, for particles of the same mass moving at the different speeds, the faster one will possess the greater kinetic energy. 12  Potential energy results from attractions of particles for each other. The particles under the influence of gravitational possess potential energy. Atoms in molecules possess potential energy because they are attracted to each other. Potential energy leads to cohesive forces in matters, which bring particles together, forming liquids and solids. Kinetic energy leads to disruptive forces in matters, which causes molecules to scatter and form gases. 13  The state of a substance depends on the relative strength of the cohesive forces (potential energy) that hold particles together and the disruptive forces (kinetic energy) that tends to scatter them. Potential energy depends on molecular size and structures and is inherent properties of the molecules; they are independent of temperature. Whereas kinetic energy is temperature dependent (molecules move faster at higher temperature). Thus, temperature plays such an important role in determining the state of matter. 14  Chapter 2 Unit conversion- significant figures- mole concept Unit conversion. Significant figures. Mole concept. Unit Conversion Is the process of converting a measurement from one unit to another while maintaining the same quantity. It involves using a conversion factor, which is a ratio that expresses how one unit relates to another. Unit conversions are commonly used in science, engineering, and everyday life to translate values between different systems of measurement (e.g., converting meters to feet, kilograms to pounds, or Celsius to Fahrenheit). Factors regulates the unit conversion Conversion Factors. Dimensional Consistency Precision and Significant Figures Unit System Scientific vs. Practical Context Historical Definitions of Units 16 Soft conversion: An exact conversion, usually within the same system or between related systems(e.g., 1 meter = 100 centimeters). Hard conversion or an adaptive conversion: An approximate or rounded conversion, often used when switching between measurement systems (e.g., 1 litre ≈ 0.26 gallons). Conversion factors: is the process of transforming a measured quantity expressed in one unit (such as grams, litres, or moles) to its equivalent in another unit using a conversion factor, which is the ratio between two equivalent values expressed in different units. Measurements Base Quantity Name of unit Symbol Length meter m Mass Kilogram Kg Time Second s Electrical current Ampere A Temperature Kelvin K Amount of substance Mole mol Luminous intensity candela cd Length 1 m — 100 cm 1 cm - 0.01 m - 10 mm 1 inch - 2.54 cm (exactly) Mass 19 Time Temperature 20 Prefix Symbol Multiple of Base Unit Giga G 1,000,000,000 or 109 Mega M 1,000,000 or 106 kilo k 1,000 or 103 deci d 0.1 or 10-1 centi c 0.01 or 10-2 milli m 0.001 or 10-3 micro m 0.000001 or 10-6 nano n 10-9 pico p 10-12 Femto f 10-15 Significant Figures and Scientific Notation The measuring device determines the number of significant figures a measurement has. Significant figures reflect the accuracy of a result or measurement. We need: To determine the correct number of significant figures (sig figs) to record in a measurement To count the number of sig figs in a recorded value To determine the number of sig figs that should be retained in a calculation. Significant Figures Figure TA 1.2 Significant figures - The number of digits in a value, often Co pyr ight © 2 001 T he McGr aw- Hill Com panies, In c. Permission req uired fo r repro duction o r disp lay. a measurement, that contribute to the degree of accuracy of the value. 22  Any digit that is not zero is significant 1.234 kg 4 significant figures. Zeros between nonzero digits are significant 606 m 3 significant figures Zeros to the left of the first nonzero digit are not significant 0.08 L 1 significant figure If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg 2 significant figures If whole numbers have zeros to the right with NO DECIMAL then all zeros to the right are NOT significant 560 mg 2 significant figures If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant 0.00420 g 3 significant figures 23  Significant Figures Any known value such as conversions (1minute = 60 seconds) have an infinite number of sig figs. Do not pay attention to them when determining sig figs! There will be many instances which fall under this category as we go through the year in chemistry! PAY ATTENTION! amples of Significan Ex t Figures How many significant figures are in the following? 3.400 4 sig figs 3004 4 sig figs 300. 3 sig figs 0.003040 4 sig figs 24 How many significant figures are in each of the following measurements? 24 mL 2 significant figures 3001 g 4 significant figures 0.0320 m3 3 significant figures 6.4 x 104 molecules 2 significant figures 560 kg 2 significant figures 25  Significant Figures Rounding Steps for Rounding to Significant Figures: Identify how many significant figures you want to round the number to. Start from the left, counting significant digits, until you reach the desired number of significant figures. Look at the next digit (the one immediately following your last significant figure): 1. If the next digit is 5 or greater, round the last significant figure up by 1. 2. If the next digit is less than 5, keep the last significant figure the same (round down). 26 Example 1: Rounding to 3 Significant Figures Let’s round 4567 to 3 significant figures: The first three significant figures are 4, 5, 6. The digit after 6 is 7, which is greater than 5, so we round the last significant figure (6) up to 7. The rounded result is 4570 (since the trailing zero is needed to maintain the value in the thousands). https://chemquiz.net/sig/ 27 28  Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers. 89.332 +1.1 one significant figure after decimal point 90.432 round off to 90.4 3.70 two significant figures after decimal point -2.9133 0.7867 round off to 0.79 Multiplication or Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x 3.6666 = 16.536366 = 16.5 3 sig figs round to 3 sig figs 6.8 ÷ 112.04 = 0.0606926 = 0.061 2 sig figs round to 2 sig figs 29  Significant Figures Addition or Subtraction  Significant Figures Multiplication or Division Let's break down the process of solving the problem using significant figures: 1.First Step (Subtraction): 1. You need to perform the operation inside the parentheses first: 5.67−2.3=3.37 2. In subtraction, the result should be rounded to the least number of decimal places from the numbers involved. Here, 5.67 has 2 decimal places, and 2.3 has 1 decimal place. So, the result 3.37 should be rounded to 1 decimal place, which gives us 3.4. 2.Second Step (Multiplication): 1. Now multiply: 3.489×3.4 2. In multiplication, the result is rounded to the fewest number of significant figures. 3.489 has 4 significant figures, and 3.4 has 2 significant figures. Therefore, the result must be rounded to 2 significant figures. 3.Perform the multiplication: 4. 3.489×3.4=11.8626 Now, round 11.8626 to 2 significant figures, which gives us 12. Scientific Notation Scientific notation is a way to express very large or very small numbers more compactly by using powers of ten. It follows the format: 602,200,000,000,000,000,000,000 6.022 x 1023 0.0000000000000000000000199 1.99 x 10-23 33 Mole (mol) The amount of a substance that contains as many elementary particles (atoms, molecules or ions), where each mole has number of 6.022 × 1023 particles. 1 mole= 6.022 × 1023 particles = Avogadro’s number Na 1 mol Al = 6.02 × 1023 atoms 1 mol CO2 = 6.02× 1023 molecules 1 mol NaCl = 6.02× 1023 Na+ ions = 6.02× 1023 Cl- ions - The number of atoms in exactly 12 g of 12C is one mole Molar Mass ( Atomic weight Aw): mass (weight) of 1 mole of atoms in grams 1 mol C atoms = 12.01 g Aw of C = 12.01* g/mol 1 mol Cl atoms = 35.45 g Aw of Cl = 35.45* g/mol 1 mol Fe atoms = 55.85 g Aw of Fe = 55.85* g/mol *( get from periodic table) Think: What is the difference between the mass and weight? Atomic Mass The mass of an atom in atomic mass units (amu) 6 Atomic number C 12.01 Atomic mass The atomic mass of elements is relative to a standard atom 12 C (6 protons, 6 neutrons) Molar Mass (Atomic weight Aw) The mass of an element atoms per one mole (g/mol) = Atomic Mass numerically Molar Mass ( Molecular weight Mw): The sum of atomic weights of 1 mol of the molecule Mw of 1 mol of H2O = 2 (Aw of H) + Aw of O = (2× 1.008) + 16 = 18.02 g/mol What are the molecular weights of the following: C2H6 N2O4 C8H18O4N2S Al2(CO3)3 MgSO4.7H2O Number of moles (n) wt ( g ) n= Mw ( g / mol ) Remember: No. of particles = No. of moles × Avogadro's number Example Methane (CH4) is the principal component of the natural gas. How many moles of methane are present in 6.07 g of CH4? Mw of CH4 = 12.01 + (4× 1.008) = 16.04 g/mol Mw = 16.04 g/mol n of CH4 = 6.07 g (CH4) × ( 1mol(CH4) ) = 0.378 mol(CH4) 16.04 g (CH4) Learning check What is the number of moles in 21.5 g CaCO3? What is the mass in grams of 0.6 mol C4H10? How many atoms of Cu are present in 35.4 g of Cu?  Percent Composition of Compounds Mass percent (weight percent) of each element in a compound. n  Aw ( x ) %x = 100 Mw 𝑛 is number of atoms of each element in the compound Percent composition Determining the Formula of a Compound: empirical formula molecular formula CH2O C6H12O6 Molecular formula = (Empirical formula)𝑥 A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance Example Calculate the mass percent of each element in ethanol (C2H5OH) ? n  Aw (x) % x = 100 M w Mass of 1 mol (molar mass) of C2H5OH = 24.02+6.048+16.00= 46.07 g/mol Mass percentof C = 2 x 12.01 g/mol x 100 = 52.14 % (4 sf) 46.07 g/mol Mass percentof H = 6 x 1.008 g/mol x 100 = 13.13 % (4 sf) 46.07 g/mol Mass percent of O = 1 x46.07 16.00 g/mol x 100 = 34.73 % g/mol (4 sf) Total mass = 52.14 +13.13 + 34.73 =100% molecular empirical H2O H2O C6H12O6 CH2O O3 O N2H4 NH2  Chapter 3 Atomic Theory Atomic Structure Early atomic theories. Dalton’s Atomic Theory. Thomson’s Theory. Rutherford’s Hypothesis. Bohr’s model. Einstein, Heisenberg and Quantum Mechanics. Spectra of hydrogen atom. Quantum Numbers for Atoms.  Atomic Theory & Atomic Structure Early Atomic Theories Democritis (400 BCE) First to propose idea of atom Atom = “a” + “tomos” = cannot be cut Based solely on logic; not supported by experiments Democritus asked: If you keep breaking matter in half, how many breaks will you have to make before you can’t break it apart any further? Democritus called the smallest possible bits of matter atoms. He had no experiments to support his theory. 41 The Ancients– B.C. Believed Aristotle's theory that everything was made up of the fundamental “elements” –Earth –Wind (air) –Fire –Water 1. Dalton’s Atomic Theory - 1808 All matter is composed of atoms which cannot be subdivided. Atoms of same element are identical (size, mass, reactivity). Atoms combine to form compounds in simple, whole # ratios. Chemical reactions involve the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. John Dalton proposed the following ideas about matter: 1.matter is made up of atoms 2.Atoms cannot be divided into smaller pieces 3.All the atoms of an element are exactly alike 4.Different elements are made of different kinds of atoms. 5. Dalton pictured an atom as a hard sphere that was the same throughout. 42 2. Thomson’s Theory In 1904, English physicist Joseph J. Thomson proposes the “plum pudding” theory of the divisible atom. His model suggested that atoms consist of a big positively- charged sphere studded with negatively charged electrons like fruit in a plum pudding. Furthermore, he put forward that the charge of the positive sphere’s charge is equal to the negative charges of the electrons. Today we call the positive charged particles protons and the negative one’s electrons. 3. Rutherford’s Hypothesis Rutherford's experiments, particularly the famous gold foil experiment, led him to conclude that atoms are mostly empty space. In this experiment, alpha particles were directed at a thin sheet of gold foil. Most of the particles passed through the foil without significant deflection, indicating that the atom is mostly empty space. An atom is mostly made up of space. The atom has a small, dense, and positively charged core at its center, which called the "nucleus.“ He proposed that almost all the mass of the atom and all of its positive charge were concentrated in this tiny nucleus. The nucleus contains protons and neutrons, while electrons orbit the nucleus at a distance. 43 4. Bohr’s model 1. Quantized Energy Levels: Bohr postulated that electrons move in specific, quantized energy levels or orbits around the nucleus. 2. Stability of Orbits: Electrons in the Bohr model are stable in their orbits. This stability is a result of the balance between the attractive force of the positively charged nucleus and the centrifugal force of the electrons' motion. Electrons do not emit energy while in these stable orbits. 3. Energy Absorption and Emission: When an electron absorbs energy, it moves to a higher energy level (further from the nucleus). Conversely, when an electron loses energy, it moves to a lower energy level (closer to the nucleus). Electrons emit or absorb energy in discrete amounts, corresponding to the energy differences between the allowed energy levels. 4. Radiation Emission: Bohr's model explained the spectral lines of hydrogen. When an electron transitions from a higher energy level to a lower one, it emits a photon of energy equal to the energy difference between the levels. This leads to the production of characteristic spectral lines. 5. Limitation to Hydrogen: The Bohr model was most successful in explaining the behavior of electrons in hydrogen atoms, where there is only one electron. It struggled to account for the behavior of more complex atoms with multiple electrons, which require the more comprehensive framework of quantum mechanics. 44 5. Einstein, Heisenberg and Quantum Mechanics 1. Quantum Mechanics: Schrödinger's model is based on the principles of quantum mechanics, a branch of physics that describes the behavior of particles at the atomic and subatomic levels. Quantum mechanics incorporates the concept of wave- particle duality, where electrons exhibit both particle-like and wave-like characteristics. 2. Electron Wave Functions: In the Electron Cloud Model, electrons are not restricted to specific orbits like in the Bohr model. Instead, electrons are described by mathematical functions called wave functions. These wave functions provide information about the probability of finding an electron in a particular region of space around the nucleus. 3. Energy Levels and Orbitals: Electrons in the Electron Cloud Model are organized into energy levels and sublevels (also known as orbitals). Each energy level corresponds to a specific energy state, and within each level, there are different types of orbitals with distinct shapes and orientations. 45 4. Uncertainty Principle: The model also adheres to Heisenberg's Uncertainty Principle, which states that it is impossible to simultaneously know both the exact position and exact momentum (or speed) of an electron. This principle underlines the probabilistic nature of quantum mechanics. 5. Electron Clouds: The model represents the three- dimensional electron distribution around the nucleus as an electron cloud or electron density map. This cloud represents the regions in space where electrons are most likely to be found, based on their probability distribution described by the wave functions. 6. Electron Spin: Schrödinger's model accounts for the concept of electron spin, which introduces an additional quantum number to describe the spin of electrons. Electrons can have one of two spin values, usually referred to as "spin- up" and "spin-down." 7.Quantum Numbers: The model uses a set of quantum numbers (principal, azimuthal, magnetic, and spin) to specify the characteristics of electrons and their positions within atoms. 46 Spectra of hydrogen atom: A hydrogen atom consists of an electron orbiting its nucleus. Spectral emission occurs when an electron transitions, or jumps, from a higher energy state to a lower energy state takes place. To distinguish the two states, the lower energy state is commonly designated as n′, and the higher energy state is designated as n. The energy of an emitted photon corresponds to the energy difference between the two states. Because the energy of each state is fixed, the energy difference between them is fixed, and the transition will always produce a photon with the same energy. The spectral lines are grouped into series according to n′. 47 Series of hydrogen atom spectra When a hydrogen atom absorbs a photon, it causes the electron to experience a transition to a higher energy level. For example, n = 1, n = 2. When a photon is emitted through a hydrogen atom, the electron undergoes a transition from a higher energy level to a lower, for example, n = 3, n = 2. During this transition from a higher level to a lower level, there is the transmission of light occurs. The quantized energy levels of the atoms, cause the spectrum to comprise wavelengths that reflect the differences in these energy levels. For example, the line at 656 nm corresponds to the transition n = 3 n = 2. The transition from the first shell to any other shell – Lyman series. The transition from the second shell to any other shell – Balmer series. The transition from the third shell to any other shell – Paschen series. The transition from the fourth shell to any other shell – Brackett series. The transition from the fifth shell to any other shell – Pfund series. 48 Quantum Numbers for Atoms There are four quantum numbers that describe the properties of an electron and the “orbital” that it occupies within an atom. What are orbitals? The main division of energy for an electron within an atom is called the energy level or shell. The energy level (shell) is subdivided into distinct areas called sublevels (subshells). The sublevels are divided into orbitals. The orbitals are where the electrons are found. Each orbital can hold no more than a pair of electrons. Four Quantum Numbers describe the following: 1. distance from the nucleus 2. shape of the orbital 3. 3-dimensional positioning of the orbital 4. direction of electron’s spin within the orbital 49 The First Quantum Number - The Principal Quantum Number Abbreviated as “n” n = 1,2,3... Represents the distance of an electron from the nucleus & the main energy level of the electron. n = number of sublevels in the energy level. n2 = number of orbitals in the energy level. 2n2 = maximum number of electrons possible in the energy level. n Energy # of # of Max # of level sublevels orbitals electrons 1 1 1 1 2 2 2 2 4 8 3 3 3 9 18 4 4 4 16 32 50 The Second Quantum Number (Types of Sublevels) The second quantum number identifies the type of sublevel that the electron occupies. There are four types of sublevels in atoms. The sublevels are identified by the letters: s, p, d, f Types of sublevels There are four types of sublevels in atoms. The sublevels are designated s, p, d, f. n Energy # of Sublevel level sublevels types 1 1 1 1s 2 2 2 2s, 2p 3 3 3 3s, 3p, 3d 4 4 4 4s, 4p, 4d, 4f “s” sublevel and “p” sublevel s orbitals p orbitals have a have a spherical dumbbell shape shape 51 The Third Quantum Number The third quantum number identifies the orbital that the electron is in. An orbital is an area within a sublevel that can hold up to two electrons An “s” sublevel has one orbital A “p” sublevel has three orbitals A “d” sublevel has five orbitals An “f” sublevel has seven orbitals 52  The Fourth Quantum Number “s” Also called the spin quantum number. Can be either +1/2 or -1/2. Specifies the direction of spin of the electron on its axis. Spins are designated up or down. Electron Spin. Opposite spins produce opposite magnetic fields. +1/2 -1/2 53 What sublevel is the first to be filled with electrons? Which sublevel listed contains orbitals of the highest energy? 54  Chapter 3 Periodic table and properties of element What is the periodic table. History of periodic table. Key to the Periodic Table. Elements of periodic table. Trends in periodic table. 56 What is the PERIODIC TABLE? Shows all known elements in the universe. Organizes the elements by chemical properties. What is the history of PERIODIC TABLE? Mendeleev 1860’s @ 60+ known elements. Father of Periodic Table (P.T.). Developed table that showed relationship between properties of elements and atomic masses. Mosley explained exceptions with discovery of Atomic Number. Modern Periodic Law: properties of the elements are a periodic function of their increasing atomic number. 57 Key to the Periodic Table Elements are organized on the table according to their atomic number. Atomic Number This refers to how many protons an atom of that element has. No two elements, have the same number of protons. Atomic Mass Atomic Mass refers to the “weight” of the atom. It is derived at by adding the number of protons with the number of neutrons. Valence Electrons Valence electrons are the electrons in the outer energy level of an atom. These are the electrons that are transferred or shared when atoms bond together. 58 Rows are called “Periods” Periods = rows From left to right What do elements in a row have in common? The same number of electron shells Every element in Period 1 (1st row) has 1 shell for its electrons (H & He) All of the elements in period 2 have two shells for their electrons. It continues like this all the way down the table 59 Columns are called “Groups” or Families Column = group = families What do elements in a group have in common? same number of valence electrons (electrons in the outer shell) They share similar characteristics with the other elements in their family. Group 1: 1 valence electron Group 2: 2 valence electrons Group 13: 3 valence electrons Group 14: 4 valence electrons Group 15: 5 valence electrons Group 16: 6 valence electrons Group 17: 7 valence electrons Group 18: 8 valence electrons except He who has 2 60 Properties of Metals Shiny. Good conductors of heat and electricity Ductile (can be stretched into thin wires) Malleable (can be pounded into thin sheets) A chemical property of metal is its reaction with water which results in corrosion. Properties of Non-Metals Poor conductors of heat and electricity Not ductile or malleable Brittle and break easily Dull Many non-metals are gases. Properties of Metalloids Solids that can be shiny or dull. Conduct heat and electricity better than non-metals but not as well as metals. They are ductile and malleable. Have properties of both metals and non-metals. 61 Region: Metals Group1: Alkalai Metals 1 valence electron Very Reactive Group 2: Alkaline Earth Metals 2 valence electrons Very reactive, but less than alkali metals 62 Groups 3 – 12: Transition Metals 1-2 valence electrons Less reactive than alkaline earth metals because they don’t give away their electrons as easily Bottom 2 row are the Lanthanide & Actinide series Lanthanide Series: Actinides Series: shiny reactive metals radioactive and unstable Most found in nature Most are man-made & not stable in nature 63 Region: Metalloids Region: Nonmetals 64 Group 17: Halogens 7 valence electrons Very reactive Nonmetals Group 18: Noble Gases ▪ 8 valence electrons (except He which only has 2) ▪ “Happy” because their outer electron shell is filled! ▪ NON REACTIVE (inert) ▪ Gases ▪ Nonmetal 65 Trends in the periodic tables ☐ The arrangement of elements in the periodic table into groups and periods with predictable properties allows us to assume certain patterns in the electron configurations. ☐ Other properties of elements follow predictable patterns in the periodic table: Atomic Radius Ionization Energy Metallic Character Electronegativity 66 Atomic Radii ✓ Atomic radius is ½ the distance between the nuclei of 2 like atoms ✓ As you move down a group, atomic radius increases. ✓ As you move left to right across a period radius decreases. As you move across the period you gain electrons but you also gain protons. More + protons hold their electrons tighter. As you move down the group you gain more electrons which are in electron orbitals further from the nucleus. The nucleus has less pull the further out you move. 67 Ionization Energy (IE) ☐ Ionization Energy is the energy required to remove 1 electron from an atom. (Make a + ion). Low IE – Easy to remove the electron High IE – Hard to remove the electron ☐ As you move down a group ionization energy decrease. ☐ As you move left to right across a period ionization energy increase Atoms on the left of the periodic table give up electrons easily to achieve a more stable electron configuration. ▪ The valence electrons in atoms lower in the group less pull from the + protons in the nucleus and are easier to remove. ▪ The Noble Gases have the highest IE of all because they are in the most stable. Low IE High IE Lower IE 68 Electronegativity ☐ Electronegativity is the tendency of an atom to attract electrons in a compound. ☐ The higher the electronegativity, the more the atom is able to attract electrons to itself. ☐ As you move down a group electronegativity decreases ☐ As you move left to right across a period electronegativity increases Electronegativity increases across a period because the number of charges on the nucleus increases. That attracts the bonding pair of electrons more strongly. As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus. 69 Metallic Character ☐ Metallic character is the tendency of an element to be shiny, silver, malleable, ductile and to react to form positive ions ☐ As you move down a group metallic character increases. ☐ As you move left to right across a period metallic character decreases. Increasing Metallic Character 70 Overview of Trends 71