Inorganic and Organic Chemistry - Introduction to Chemistry PDF

INORGANIC AND ORGANIC CHEMISTRY MODULE 1: INTRODUCTION TO CHEMISTRY CHEMISTRY: THE STUDY OF A SCIENCE Empedocles and Democritus – FOR THE TWENTY-FIRST CENTURY contributed philosophical ideas about the...

INORGANIC AND ORGANIC CHEMISTRY MODULE 1: INTRODUCTION TO CHEMISTRY CHEMISTRY: THE STUDY OF A SCIENCE Empedocles and Democritus – FOR THE TWENTY-FIRST CENTURY contributed philosophical ideas about the natures of matter, anticipating some CHEMISTRY concepts later developed in chemistry.  Study that deals with the composition CLASSICAL GREECE structure, properties, and changes of matter saw limited progress in formal chemistry,  Often called “central science” but the emphasis on natural philosophy  Derived from word “kheem” (black laid the groundwork for later scientific land) inquiry, influencing the scientific o Greek word: Khemia methodology of thinkers like Aristotle. o Arabic word: Al-kimiya o English word: Alchemy HELLENISTIC PERIOD “Chemistry is everywhere” witnessed advancements in alchemy, as scholars in Alexandria and other  Composition of matter Hellenistic centers explored practical  Chemical reactions techniques, experimented with various  Materials and engineering substances, and contributed to the  Pharmaceutical and medicine evolving understanding of chemical  Food chemistry processes  Cosmetics and personal care products ISLAMIC GOLDEN AGE (8TH – 4TH MATTER CENTURIES) Anything that occupies space and has mass significant advancements in chemistry, Key characteristics: including the refinement of distillation  Mass – amount of material in an object techniques, the discovery of (kg or g) numerous chemical substances, and  Volume – space occupied by matter the development of scientific (cubic units of m or cm) methods, laid the foundation for modern chemistry. Composed of tiny particles: Atom MEDIEVAL PERIOD  basic building blocks of matter  smallest unit of an elementthat retain the Alchemy (mystical and proto-scientific chemical properties of that element. practice) flourished with alchemists Molecules seeking to transmute base metals into  Groups of two or more atoms that are gold and discover the “elixir of life”. chemically bonded together RENAISSANCE PERIOD HISTORICAL DEVELOPMENT Marked a transition in chemistry as alchemical practices merged with PRE-HISTORIC, ANCIENT, AND CLASSICAL emerging empirical methods setting the GREECE AND HELLENISTIC PERIOD PRE- stage for the scientific revolution, with HISTORIC GREEK figures like Societies laid rudimentary foundations Paracelsus contributing to the for chemistry through practical understanding of chemical principles and knowledge of metallurgy, pottery, and medicinal applications. early medicinal practices using herbs and minerals. AGE OF ENLIGHTENMENT ANCIENT GREECE INORGANIC AND ORGANIC CHEMISTRY MODULE 1: INTRODUCTION TO CHEMISTRY Chemists like Antoine Lavoisier Deals with compounds that do not revolutionized the field by establishing contain carbon-hydrogen (C-H) the principles of modern chemistry, bonds. including the law of conservation of Includes the study of metals, minerals, mass and the identification of salts, and coordination compounds elements RISE OF MODERN CHEMISTRY ORGANIC CHEMISTRY Focuses on the study of carbon  Began with Lavoisier’s foundational containing compounds. contributions in the late 18th century, Abundant in living organisms. followed by Mendeleev’s periodic law and advancements in industrial PHYSICAL CHEMISTRY chemistry in the 19th century, culminating in the 20th century with the Combines principles of physics and integration of quantum mechanics and chemistry to study the physical properties the intersection with molecular biology, and behavior of matter. shaping chemistry into a comprehensive multidisciplinary science. ANALYTICAL CHEMISTRY ADVANCES IN MODERN CHEMISTRY Involves the identification and quantification of chemical substances Encompasses a wide array of achievements, including the BIOCHEMISTRY development of new materials, Explores the chemical processes within breakthroughs in nanotechnology, the and related to living organisms. understanding of complex biological processes, and the design of novel ENVIRONMENTAL CHEMISTRY pharmaceuticals. Focuses on the study of chemical BRANCHES OF CHEMISTRY processes occurring in the environment and their impact on ecosystems NUCLEAR CHEMISTRY Involves the study of radioactive materials and nuclear reactions. MATERIALS CHEMISTRY Examines the synthesis, structure, properties, and applications of materials. POLYMER CHEMISTRY Focuses on the study of polymers. Large molecules composed of repeating structural units. CONCLUSION Chemistry is often referred to as a central science. INORGANIC CHEMISTRY INORGANIC AND ORGANIC CHEMISTRY MODULE 1: INTRODUCTION TO CHEMISTRY The behavior of atoms, molecules, and ions determines the sort of world we live in, our shapes and sizes, and even how we feel on a given day. Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER MODULE 2: MATTER- ITS STATES, CLASSIFICATION, CHANGES & TRANSITION.  MATTER: anything that has mass and occupies space  MASS refers to the amount of matter present in a sample.  Matter can be changed into energy and energy into matter. o Energy as the ability to do work, or to move matter. o All forms of energy exhibit both Kinetic (moving) and Potential (stored). FORMS OF ENERGY  CHEMICAL ENERGY – is stored in the bonds of a chemical substances.  ELECTRICAL ENERGY – results from the movement of charged particles.  MECHANICAL ENERGY – directly involved in moving matter.  RADIANT ENERGY - is the form of energy associated with the movement of light, electromagnetic waves or particles.  NUCLEAR ENERGY- is the energy released when the nuclei of atoms are split or fused.  THERMAL ENERGY- is heat energy, or the energy of moving or vibrating molecules. STATES OF MATTER (a) A solid has a definite shape and a definite volume. (b) A liquid has an indefinite shape—it takes the shape of its container—and a definite volume. (c) A gas has an indefinite shape and an indefinite volume—it assumes the shape and volume of its container.  Plasma state is the 4th state of matter. It has hot ionized gas, a gas into which sufficient energy is provided to free electrons from atoms or molecules and to allow both species i.e. ions or electrons to co-exist. Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER  BEC (Bose-Einstein Condensate) is the 5th state of matter of a dilute gas of bosons cooled to a temperature very close to absolute zero.  Under such conditions, a large fraction of bosons occupy the lowest quantum state, at which point macroscopic quantum phenomena become apparent.  For example, Helium, they are amazingly cold, too close to 270∘C or zero kelvin.  The main difference between plasma and Bose-Einstein condensate is that the plasma state contains a gas of ions and free electrons, whereas Bose-Einstein condensate contains a gas of bosons at low densities, which is cooled to a low temperature close to absolute zero. SUMMARY TABLE FOR STATES OF MATTER PROPERTIES OF MATTER  A property is a distinguishing characteristic of a substance that is used in its identification and description Two categories: A. PHYSICAL PROPERTY  is a characteristic of a substance that can be observed without changing the basic identity of the substance. Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER  Common physical properties include color, odor, physical state (solid, liquid, or gas), melting point, boiling point, and hardness  During the process of determining a physical property, the physical appearance of a substance may change, but the substance’s identity does not. B. CHEMICAL PROPERTY  A chemical property is a characteristic of a substance that describes the way the substance undergoes, or resists change to form a new substance.  For example, copper objects turn green when exposed to moist air for long periods of time, this is a chemical property of copper. The green coating formed on the copper is a new substance that results from the copper’s reaction with oxygen, carbon dioxide, and water present in air.  The properties of this new substance (the green coating) are very different from those of metallic copper. On the other hand, gold objects resist change when exposed to air for long periods of time. The lack of reactivity of gold with air is a chemical property of gold. SUBCLASSIFICATION Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER CHANGES IN MATTER  A physical change is a process in which a substance changes its physical appearance but not its chemical composition.  A chemical change is a process in which a substance undergoes a change in chemical composition Describe what change/s is/are happening with the scenario Answer: This illustration depicts both physical and chemical change. Melting of the wax in an example physical change. The wax can be molded again to form a new candle. However, burning the wick of the candle is an example of chemical change. The white wick will eventually burn and become black until consumed which is irreversible. CLASSIFICATION OF MATTER Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER  A pure substance is a single kind of matter that cannot be separated into other kinds of matter by any physical means.  All samples of a pure substance contain only that substance and nothing else.  A pure substance always has a definite and constant composition  A mixture is a physical combination of two or more pure substances in which each substance retains its own chemical identity.  Components of a mixture retain their identity because they are physically mixed rather than chemically combined.  Consider a mixture of small rock salt crystals and ordinary sand. Mixing these two substances changes neither the salt nor the sand in any way. The larger, colorless salt particles are easily distinguished from the smaller, light-gray sand granules  A heterogeneous mixture is a mixture that contains visibly different phases (parts), each of which has different properties.  A nonuniform appearance is a characteristic of all heterogeneous mixtures.  Examples include chocolate chip cookies and blueberry muffins.  Naturally occurring heterogeneous mixtures include rocks, soils, and wood  A homogeneous mixture is a mixture that contains only one visibly distinct phase (part), which has uniform properties throughout.  The components present in a homogeneous mixture cannot be visually distinguished.  A sugar–water mixture in which all of the sugar has dissolved has an appearance similar to that of pure water. Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER  Air is a homogeneous mixture of gases; motor oil and gasoline are multicomponent homogeneous mixtures of liquids; and metal alloys such as 14-karat gold (a mixture of copper and gold) are examples of homogeneous mixtures of solids. TRANSITION OF MATTER  happens when a substance changes from a solid, liquid, or gas state to a different state. Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER SUMMARY  Matter anything that has mass and occupies space  Mass refers to the amount of matter present in a sample.  Matter can be changed into energy and energy into matter. Forms of energy include Chemical, Electrical, Mechanical, Radiant, Thermal & Nuclear.  Commonly, three physical states exist for matter: solid, liquid, and gas. But as science progresses, additional states were discovered such as plasma and Bose-Einstein condensate. Properties of matter are of two general types: physical and chemical Physical properties of matter may be classified as to INTENSIVE or EXTENSIVE property  Changes in matter can happen in 2 manners, physically and chemically.  In addition to its classification by physical state, matter can also be classified in terms of its chemical composition as a pure substance or as a mixture Inorganic and Organic Chemistry (Lecture) MODULE 2 - MATTER  Phase transition happens when a substance changes from a solid, liquid, or gas state to a different state.  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY MEASUREMENTS IN CHEMISTRY MEASUREMENT SYSTEMS  What is a Measurement? o Measurement is the determination of the dimensions, capacity, quantity, or extent of something.  Systems of Measurements o English System (commerce):  inch, foot, pound, quart, and gallon o Metric System (scientific work):  gram, meter, and liter  More convenient to use (decimal unit system). METRIC SYSTEM UNITS  There is one base unit for each type of Metric Mass Units measurement (length, mass, volume, etc.).  Gram (g):  Add prefixes to the base unit to indicate o Base unit of mass the size of the unit.  Mass is measured by determining the  The prefix is independent of the base unit amount of matter in an object. and always remains constant. o Mass – measure of the total quantity of matter in an object o Weight – measure of the force Common Metric System Prefixes exerted on an object by gravitational forces Metric Volume Units Metric Length Units  Liter (L): o Base unit of volume  Meter (m):  Volume is measured by determining the o Base unit of length. amount of space occupied by a three-  Length is measured by determining the dimensional object. distance between two points. o 1 liter = 1000 cm3 = 1 dm3  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY o 1 mL = 1 cm 3  Results any time a measurement is o mL generally used for liquids and made. gases. o cm3 used for solids UNCERTAINTY IN MEASUREMENT AND SIGNIFICANT CHANGES Uncertainty in Measurements  A digit that must be estimated is called uncertain.  A measurement always has some degree of uncertainty.  Record the certain digits and the first uncertain digit (the estimated number). Consider These Rulers Measurements made with ruler A will have greater uncertainty than those made with ruler B. A cube 10 cm on a side has a volume of 1000 cm3, which is equal to 1 L. A cube 1 cm on a side has a volume of 1 cm3, which is equal to 1 mL. EXACT AND INEXACT NUMBERS Exact Numbers  A number whose value has no Ruler B is more precise than Ruler A. uncertainty associated with it – that is, it is known exactly. Significant Figures o Definitions – 12 objects in a dozen  Digits in a measurement that are known o Counting – 15 pretzels in a bowl with certainty plus one digit that is o Simple fractions – ½ or ¾ estimated. Inexact Numbers  A number whose value has a degree of Guidelines for Determining Significant uncertainty associated with it. Figures  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY  In any measurement, all nonzero digits 1. In multiplication and division, the are significant. number of significant figures in the o 3456 has 4 significant figures. answer is the same as the number of  Zeros may or may not be significant significant figures in the measurement because zeros can be used in two ways: that contains the fewest significant o to position a decimal point figures. o to indicate a measured value.  Zeros that perform the first function are not significant, and zeros that perform the second function are significant. When zeros are present in a measured number, we follow these rules: ▪ Leading zeros are zeros that are at the beginning of a number. These do not count as significant figures. o 0.048 has 2 significant figures. ▪ Confined zeros are zeros between nonzero digits. These always count as 2. In addition, and subtraction, the significant figures. answer has no more digits to the right of o 16.07 has 4 significant figures. the decimal point than are found in the ▪ Trailing zeros are zeros at the right end measurement with the fewest digits to the of the number. They are significant only if right of the decimal point. the number contains a decimal point. o 9.300 has 4 significant figures ▪ Trailing zeros are zeros at the right end of the number. They are not significant if the number lacks an explicitly shown decimal point. o 150 has 2 significant figures. SIGNIFICANT FIGURES AND MATHEMATICAL OPERATIONS Concept Check Rounding Off Numbers You have water in each graduated cylinder shown. You then add both samples to a beaker. ▪ Process of deleting unwanted (nonsignificant) digits from calculated How would you write the number describing the numbers. total volume? 1. If the first digit to be deleted is 4 or less, 3.1 mL simply drop it and all the following digits. What limits the precision of this number? ▪ 5.83298 becomes 5.83 (for 3 sig figs). 2. If the first digit to be deleted is 5 or greater, that digit and all that follow are dropped, and the last retained digit is increased by one. ▪ 7.86541 becomes 7.87 (for 3 sig figs). Operational Rules  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY Converting from Decimal to Scientific Notation ▪ The decimal point in the decimal number is moved to the position behind (to the right of) the first nonzero digit. ▪ The exponent for the exponential term is equal to the number of places the decimal point has moved. o 300. written as 3.00 × 102 (three  The total volume is 3.1 mL. significant figures)  The first graduated cylinder limits the o 0.004890 written as 4.890 × 10–3 precision of this number with a volume of (four significant figures) 2.8 mL.  The second graduated cylinder has a volume of 0.28 mL. Therefore, the final volume must be 3.1 mL since the first volume is limited to the tenths place. Exact Numbers ▪ Because exact numbers have no uncertainty associated with them, they possess an unlimited number of significant figures. o 1 inch = 2.54 cm, exactly. o 9 pencils (obtained by counting). ▪ Exact numbers never limit the The exponent is positive if the original decimal number of significant figures in a number is 10 or greater and is negative if the computational answer. original decimal number is less than 1. For numbers between 1 and 10, the exponent is zero. SCIENTIFIC NOTATION Multiplication and Division in Scientific Notation Exponential Notation 1. To multiply exponential terms, add the ▪ A numerical system in which numbers exponents. are expressed in the form A × 10n where 2. To divide exponential terms, subtract A is a number with a single nonzero digit the exponents. to the left of the decimal place and n is a whole number. CONVERSION FACTORS o A is the coefficient o n is a whole number ▪ A ratio that specifies how one unit of measurement is related to another unit of measurement.  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY ▪ To convert from one unit to another, use ▪ Identify the known or given quantity (both the equivalence statement that relates numerical value and units) and the units the two units. of the new quantity to be determined. o 1 ft = 12 in. o 6.8 ft =? in. ▪ The two conversion factors are: ▪ Multiply the given quantity by one or more conversion factors in such a manner that 1 ft 12 in. the unwanted (original) units are and ▪ canceled, leaving only the desired units. The two conversion factors are: 12 in. 1 ft Equalities and Conversion Factors That 1 ft 12 in and Relate the English and Metric Systems of 12 in 1 ft Measurement ▪ Perform the mathematical operations indicated by the conversion factor setup. Example #2 ▪ An iron sample has a mass of 4.50 lb. What is the mass of this sample in grams? ▪ (1 kg = 2.2046 lbs; 1 kg = 1000 g) DIMENSIONAL ANALYSIS Steps for Using Dimensional Analysis ▪ Use when converting a given result from one system of units to another: 1. Identify the known or given DENSITY quantity (both numerical value and units) and the units of the ▪ Density is the ratio of the mass of an object to the volume occupied by that new quantity to be determined. 2. Multiply the given quantity by object one or more conversion factors ▪ The equation for density is: in such a manner that the unwanted (original) units are mass canceled, leaving only the Density = desired units. volume 3. Perform the mathematical operations indicated by the ▪ generally expressed in grams per cubic conversion factor setup. centimeter (g/cm3) for solids, grams per milliliter (g/mL) for liquids, and grams per Example #1 liter (g/L) for gases. ▪ A golfer putted a golf ball 6.8 ft across Because density is a property that does not a green. How many inches does this depend on the quantity of mass present, for a represent? given substance the ratio of mass to volume always remains the same. In other words, V  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY increases as m does. Density usually decreases ▪ This use of density enables us to with temperature calculate the volume of a substance if we know its mass. ▪ Because gas densities are often very low, we express them in units of grams per ▪ Conversely, the mass can be calculated liter (g/L): if the volume is known. 1 g/cm3 = 1 g/mL = 1000 kg/m3 ▪ Density conversion factors, like all other conversion factors, have two reciprocal 1 g/L = 0.001 g/Ml forms. ▪ For a density of 1.03 g/mL, the two conversion factor forms are Example #1 ▪ A certain mineral has a mass of 17.8 g and a volume of 2.35 cm3. What is the density of this mineral? Example #2 ▪ What is the mass of a 49.6 mL sample of a liquid, which has a density of 0.85 g/mL? Density as a Conversion Factor ▪ Density can be used as a conversion factor that relates the volume of a substance to its mass. TEMPERATURE SCALES  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY Three Systems for Measuring Temperature for the freezing point of water and 212 for the normal boiling point of water. ▪ Celsius - divides the range between the freezing point (0°C) and boiling point Besides the boiling and freezing point for water, a (100°C) of water into 100 degrees. third reference point is shown in Figure 2.10 for each of the temperature scales—normal human ▪ Kelvin - SI base unit of temperature: It is body temperature the absolute temperature scale Conversions Between Temperature Scales ▪ Fahrenheit - defines the normal freezing and boiling points of water to be exactly ▪ Because the size of the degree is the 32°F and 212°F, respectively. same, the relationship between the Kelvin and Celsius scales is very simple. The Three Major Temperature Scales No conversion factors are needed; all that is required is an adjustment for the differing numerical scale values. ▪ The adjustment factor is 273, the number of degrees by which the two scales are offset from one another. The Celsius scale is the scale most commonly ▪ The relationship between the Fahrenheit encountered in scientific work. The normal boiling and Celsius scales can also be stated in and freezing points of water serve as reference an equation format. points on this scale, the former having a value of 100 and the latter 0. Thus, there are 100 “degree intervals” between the two reference points The Kelvin scale is a close relative of the Celsius scale. Both have the same size degree, and the number of degrees between the freezing and boiling points of water is the same. The two scales differ only in the numbers assigned to the reference points. On the Kelvin scale, the boiling point of water is 373 kelvins (K) and the freezing point of water is 273 K. The choice of these reference points makes all temperature readings on the Kelvin scale positive values. Note that the degree sign () is not used with the Kelvin scale. For example, we say that an object has a temperature of 350 K (not 350K). The Fahrenheit scale has a smaller degree size than the other two temperature scales. On this scale, there are 180 degrees between the Exercise: freezing and boiling points of water as contrasted to 100 degrees on the other two scales. Thus, the ▪ At what temperature does C = F? Celsius (and Kelvin) degree size is almost two Solution: times (9–5) larger than the Fahrenheit degree. Reference points on the Fahrenheit scale are 32  INORGANIC AND ORGANIC CHEM LEC – MODULE 3: MEASUREMENTS IN CHEMISTRY ▪ Since °C equals °F, they both should be results of three successive weighing by the same value (designated as variable each student are x). ▪ Use one of the conversion equations such as: 5 C   F  32  The true mass of the wire is 2.000 g. Therefore, Student B’s results are more precise than those of Student A (1.970 g, 1.972 g, and 1.968 g 9 deviate less from 1.970 g than 1.964 g, 1.971 g, and 1.978 g from 1.971 g), but neither set of ▪ Substitute in the value of x for both °C results is very accurate. Student C’s results are and °F. Solve for x. not only the most precise, but also the most accurate, because the average value is closest to the true value. Highly accurate measurements are usually precise, too. On the other hand, highly precise measurements do not necessarily guarantee accurate results. For example, an improperly calibrated meterstick or a faulty balance may give precise readings that are in error ▪ The terms accuracy and precision are related to two types of error encountered in measurements. ▪ Systematic errors occur in a ACCURACY AND PRECISION predictable manner, leading to measured values that are ▪ Accuracy tells us how close a constantly off from the true value. measurement is to the true value of the ▪ Random errors are not quantity that was measured. predictable, leading to measured ▪ Precision refers to how closely two or values that vary greatly from the more measurements of the same true value. quantity agree with one another ▪ Suppose, for example, that three students are asked to determine the mass of a piece of copper wire. The  INORGANIC AND ORGANIC CHEMISTRY LEC – ATOMS AND ITS PROPONENTS 1 MODULE 4 – ATOMS AND ITS PROPONENTS  He disproved Thomson’s model thru “Gold Foil/ Film Experiment” HISTORY AND PROPONENTS OF ATOMIC  He concluded the following: MODELS o atom is just an empty space, the Democritus nucleus accounts for the (+) charge & mass of the atom & electrons  expressed the belief that all matter consists scattered around the nucleus of very small, indivisible particles, which he “Nuclear Model” named atomos (meaning uncuttable or indivisible) Neils Bohr John Dalton  “Planetary Model”  formulated a precise definition of the  The atom consists of nucleus surrounded by indivisible building blocks of matter that we electrons traveling in circular orbits called called atoms orbitals Joseph John Thomson Erwin Schrödinger  He said that atom is a sphere of (+) particles  “Quantum Mechanical Model” to which are embedded (-) particles  Electron moves in 3D space (electron cloud)  Raisin- Bread Model or Plum- Pudding  Modern atomic structure Model Ernest Rutherford INORGANIC AND ORGANIC CHEMISTRY LEC – ATOMS AND ITS PROPONENTS 2 DALTON’S ATOMIC THEORY Protons and electrons carry the same amount of charge; the charges, however, are opposite (positive 1. Elements are composed of extremely small versus negative). particles, called atoms. 2. All atoms of a given element are identical, Both protons and neutrons are massive particles having the same size, mass, and chemical compared to electrons; they are almost 2000 times properties. The atoms of one element are heavier. different from the atoms of all other elements. 3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements presents is either an integer or a simple fraction. 4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. SUBATOMIC PARTICLES  Subatomic particle is a very small particle ATOMIC NUMBER AND MASS NUMBER that is a building block for atoms Atomic Number  Electron o Negatively charge particle  The number of protons in the nucleus of o It is the smallest, in terms of mass, of each atom of an element the three types of subatomic particles.  Proton Mass Number o Positively charge particle  Neutron  The total number of neutrons and protons o No charge particle present in the nucleus of an atom of an o Neutral element  Nucleus o a dense central core within the atom All atoms can be identified by the number of protons and neutrons they contain.  To maintain electrical neutrality, an atom must contain an equal number of positive and negative charges.  INORGANIC AND ORGANIC CHEMISTRY LEC – ATOMS AND ITS PROPONENTS 3 Determine the number of neutrons present. Determine the number of electrons present. Answers: a. 11 protons; b. 12 neutrons; c. 11 electrons Compute what is asked and fill in the table with the correct mass number, number of electrons, neutrons, and protons. # of protons = atomic number # of electrons = atomic number # of neutrons = mass number – atomic number ISOTOPES, ISOBARS AND ISOTONES Isotopes  atoms of an element that have the same number of protons and the same number of electrons but different numbers of neutrons Isotopes of an element have the same chemical properties, but their physical properties are often slightly different. Isotopes of an element have the same chemical properties because they have the same number of electrons. They have slightly different physical properties because they have different numbers of neutrons and therefore different masses. PRACTICE EXERCISE: Hydrogen-1 is usually called hydrogen but is An atom has an atomic number of 11 and a occasionally called protium. Hydrogen-2 has the name mass number of 23. deuterium (symbol D), and hydrogen-3 is called tritium (symbol T). Determine the number of protons present. Isobars  INORGANIC AND ORGANIC CHEMISTRY LEC – ATOMS AND ITS PROPONENTS 4  different elements have the same atomic weight or mass number but different atomic number Isotones  different elements having the same number of neutrons QUANTUM NUMBERS Three numbers are derived from the mathematical solution of the Schrödinger equation for the hydrogen atom:  the principal quantum number (n)  the angular momentum quantum number (l),  the magnetic quantum number (ml) 1. The principal quantum number (n) n = 1, 2, 3, 4, …. The value of n determines the energy of an orbital and relates to the average distance of the electron from the nucleus in a particular orbital The larger n is, the greater the average distance of an electron in the orbital from the nucleus and therefore the larger the orbital.  INORGANIC AND ORGANIC CHEMISTRY LEC – ATOMS AND ITS PROPONENTS 5 2. The angular momentum quantum number (l) for The magnetic quantum number (m/) a given value of n, l = 0, 1, 2, 3, … n-1 describes the orientation of the orbital in space. Shape of the “volume” of space that the e- occupies (orbitals) 4. The electron spin quantum number (ms) Experiments on the emission spectra of hydrogen and sodium atoms indicated that 3. The magnetic quantum number (ml) lines in the emission spectra could be split by the application of an external magnetic field.  INORGANIC AND ORGANIC CHEMISTRY LEC – ATOMS AND ITS PROPONENTS 6 directed through a magnetic field. For example, when a hydrogen atom with a single electron passes through the field, it is deflected in one direction or the other, depending on the direction of the spin. In a stream consisting of many atoms, there will be equal distributions of the two kinds of spins, so that two spots of equal intensity are detected on the screen. Pauli exclusion principle Four quantum numbers characterize each electron no two electrons in an atom can have the in an atom: same four quantum numbers. the principal quantum number n identifies the main energy level, or shell, of the orbital the angular momentum quantum number ℓ indicates the shape of the orbital; the magnetic quantum number mℓ specifies the orientation of the orbital in space the electron spin quantum number ms indicates the direction of the electron’s spin on its own axis. The (a) clockwise and (b) counterclockwise spins of an electron. The magnetic fields generated by these two spinning motions are analogous to those from the two magnets. The upward and downward arrows are used to denote the direction of spin Experimental arrangement for demonstrating the spinning motion of electrons. A beam of atoms is  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS HISTORY OF THE PERIODIC TABLE  Proposed a system of chemical symbols based on the first letter of their name and History some based on their Latin name.  During the mid-nineteenth century, Antoine Lavoisier scientists began to look for order in the increasing amount of chemical  Arranged elements into groups of simple information that had become available. substances that do not decompose by  Scientists knew that certain elements any means. had properties that were very similar to those of other elements.  So, they sought reasons for these similarities in the hope that these similarities would suggest a method for arranging or classifying the elements. Periodic Table  A period table is a tabular arrangement of the elements in order of increasing Alexandre-Émile Béguyer de Chancourtois atomic number such that elements having similar chemical properties are  Made a list of elements arranged by positioned in vertical columns. increasing atomic weight in spiral order. The Periodic Law  This law states that when elements are arranged in order of increasing atomic number, elements with similar chemical properties occur at periodic intervals. Note!  The term "Periodic" means regularly occurring Johann Dobereiner  Before establishing the modern periodic table that we know, many scientists in the  He classified core elements into groups past have attempted to organize the of three, which he called triads. elements.  The elements in a triad had similar chemical properties and orderly physical Jons Jakob Berzelius properties, John Newlands  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS  He suggested that elements be arranged Period in octaves because he noticed (after arranging the elements in order of  A period is a horizontal row of elements increasing atomic mass) that certain in the periodic table. properties repeated after every 8th Group element  A vertical column of elements in the Loathar Meyer & Dmitri Mendeleev periodic table. THE PERIODIC TABLE  Observe the arrangement and names of the different groups and periods.  Notice how certain elements with similar  Arranged the elements according to its chemical properties are grouped atomic weights (mass) together. Henry Moseley  The groups and periods give the location of the elements.  For example, Iodine is Element 53.  A British chemist responsible for the  Iodine is in period 5, group VIIA arrangement of the periodic table in or 17 terms of atomic number. GROUPS AND PERIODS OF ELEMENTS  The location of an element within the periodic table is specified by giving its period number and group number.  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS What is the noble gas located in period 4?  Answer: Kr (Krypton) THE 18 GROUPS IN THE PERIODIC TABLE GROUPS IN THE PERIODIC TABLE  There are 18 groups in the periodic table.  Group A Elements - Representative Elements  Group B Elements - Transition  The elements in the periodic table have Elements their specifications indicated. The Group A Elements (Representative Elements) have other names. PROPERTIES OF THE GROUPS  Alkali Metals - A general name for any element in Group IA, excluding Hydrogen. What is the element located in both Period 3 and  These are soft, shiny metals that Group IVA? readily react with water.  Answer: Si (Silicon)  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS  Alkaline Earth Metals - A general name for any element in Group IIA.  These are soft, shiny metals, but they are only moderately reactive with water.  Halogens - is a general name for any element in Group VIIA.  These are reactive elements that are gases at room temp. or become such at temperatures slightly above room temp.  Noble Gases - These are on the extreme right of the periodic table. Group VIIIA.  These are unreactive gases that undergo few, if any, chemical reactions. Provide the symbol, period and group of:  Neon  Answer: Ne, Period 2, Group VIIIA Provide the symbol, period and group of:  Fluorine  Answer: F, Period 2, Group VIIA Provide the symbol, period and group of:  Lead  Answer: Pb, Period 6, Group IVA THE SHAPE OF THE PERIODIC TABLE  The periodic table is arranged by increasing atomic number.  However, this pattern is violated in Groups IIIB and IVB.  Element 57 follows element 72.  Element 89 follows element 104.  The missing elements (Elements 58 to 71 and Elements 90 to 103) are located in the two rows at the METALS, NONMETALS, AND METALLOIDS bottom of the periodic table.  This was done to make a more compact  Observe the location of the Metals, periodic table. Nonmetals and Metalloids. CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS  Chemically, they behave mostly as nonmetals. METALS  It is an element with luster, thermal conductivity, electrical conductivity, and THE PERIODIC TRENDS malleability.  Except for mercury, all metals are solids at room temperature.  Metals are good conductors of heat and electricity.  Most metals are ductile, can be drawn into wires, are malleable, and can be rolled into sheets.  Most metals have high luster or shine, high density, and high melting points.  Among the more familiar metals are the elements iron, aluminum, copper, silver, gold, lead thin, and zinc ATOMIC RADIUS NONMETALS  The measure of the size of the element's atoms  It is an element without luster, thermal  The distance from the nucleus to the conductivity, electrical conductivity, and boundary of the surrounding cloud of malleability. electrons  Many nonmetals, such as hydrogen, oxygen, nitrogen, and noble gases, are The atomic radius is one-half the distance gases. between the nuclei of two atoms (just like a radius  The only nonmetal that is a liquid at room is half the diameter of a circle). However, this idea temperature is bromine. is complicated by the fact that not all atoms are  Solid nonmetals include carbon, iodine, normally bound together in the same way. Some sulfur, and phosphorus. are bound by covalent bonds in molecules, some  Nonmetals have lower densities and are attracted to each other in ionic crystals, and lower melting points than metals. others are held in metallic crystals. Nevertheless, it is possible for a vast majority of elements to METALLOID form covalent molecules in which two like atoms are held together by a single covalent bond. The  It is an element with a metallic covalent radii of these molecules are often appearance. referred to as atomic radii. This distance is  These are brittle and only fair conductors measured in picometers. Atomic radius patterns of electricity. are observed throughout the periodic table.  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS Atomic size gradually decreases from left to right Compare Aluminum and Fluorine based on across a period of elements. This is because, electronegativity. within a period or family of elements, all electrons are added to the same shell. However, at the  ANSWER: Fluorine has higher same time, protons are being added to the electronegativity nucleus, making it more positively charged. The effect of increasing proton number is greater than that of the increasing electron number; therefore, there is a greater nuclear attraction. This means that the nucleus attracts the electrons more strongly, pulling the atom's shell closer to the nucleus. The valence electrons are held closer towards the nucleus of the atom. As a result, the atomic radius decreases. Down a group, atomic radius increases. The valence electrons occupy higher levels due to the increasing quantum number (n). As a result, the valence electrons are further away from the nucleus as ‘n’ increases. Electron shielding prevents these outer electrons from being attracted to the nucleus; thus, they are loosely held, and the resulting atomic radius is large. IONIZATION ENERGY TRENDS:  Energy required to remove electrons  Atomic radius decreases from left to right from gaseous atoms or ions within a period. This is caused by  Measure of the tendency of an atom or the increase in the number of protons ion to surrender an electron or the and electrons across a period. One strength of electron binding proton has a greater effect than one  Also called ionization potential electron; thus, electrons are pulled  Measured in volts towards the nucleus, resulting in a smaller radius.  Atomic radius increases from top to bottom within a group. This is caused by electron shielding. Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous phase. Conceptually, ionization energy is the opposite of electronegativity. The lower this energy is, the more readily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely it is the atom becomes a cation.  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS Generally, elements on the right side of the than the atom above it (this is the atomic radius periodic table have a higher ionization energy trend, discussed below). This means that an because their valence shell is nearly filled. added electron is further away from the atom's Elements on the left side of the periodic table nucleus compared with its position in the smaller have low ionization energies because of their atom. With a larger distance between the willingness to lose electrons and become cations. negatively charged electron and the positively- Thus, ionization energy increases from left to charged nucleus, the force of attraction is right on the periodic table. relatively weaker. Therefore, electron affinity decreases. Moving from left to right across a TRENDS: period, atoms become smaller as the forces of  The ionization energy of the elements attraction become stronger. This causes the within a period generally increases from electron to move closer to the nucleus, thus left to right. This is due to valence shell increasing the electron affinity from left to right stability. across a period.  The ionization energy of the elements within a group generally decreases from TRENDS: top to bottom. This is due to electron  Electron affinity increases from left to shielding. right within a period. This is caused by  The noble gases possess very high the decrease in atomic radius. ionization energies because of their full valence shells as indicated in the graph.  Electron affinity decreases from top to Note that helium has the highest bottom within a group. This is caused by ionization energy of all the elements. the increase in atomic radius. ELECTRON AFFINITY ELECTRONEGATIVITY  Change in energy in kJ/mole of a neutral  The tendency of an atom to attract a atom (in the gaseous phase) when an shared pair of electrons or electron electron is added to the atom to form a density towards itself negative ion  It is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus  Compare Lithium and Potassium in terms of atomic radius  Answer: Potassium has higher atomic radius  Compare Aluminum and Fluorine based on electronegativity.  Answer: Fluorine has higher electronegativity Electron affinity is the ability of an atom to accept an electron. Unlike electronegativity, electron Electronegativity measures an atom's tendency affinity is a quantitative measurement of the to attract and form bonds with electrons. This energy change that occurs when an electron is property exists due to the electronic configuration added to a neutral gas atom. The more negative of atoms. Most atoms follow the octet rule (having the electron affinity value, the higher an atom's the valence, or outer, shell comprise of 8 affinity for electrons. electrons). Because elements on the left side of the periodic table have less than a half-full Electron affinity generally decreases down a valence shell, the energy required to gain group of elements because each atom is larger electrons is significantly higher compared with the  CHEML-MODULE 5 – PERIODIC TABLE AND ITS 09/10/2024 ELEMENTS energy required to lose electrons. As a result, the ry/Periodic_Trends_of_Elemental_Properties/Pe elements on the left side of the periodic table riodic_Trends generally lose electrons when forming bonds. Conversely, elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a complete valence shell of 8 electrons. The nature of electronegativity is effectively described thus: the more inclined an atom is to gain electrons, the more likely that atom will pull electrons toward itself. From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one. From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius. Important exceptions of the above rules include the noble gases, lanthanides, and actinides The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values. As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements. According to these two general trends, the most electronegative element is fluorine, with 3.98 Pauling units REFERENCE ON ADDITIONAL INFORMATIONS: https://chem.libretexts.org/Bookshelves/Inorgani c_Chemistry/Supplemental_Modules_and_Webs ites_(Inorganic_Chemistry)/Descriptive_Chemist