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Introduction to General Chemistry

Introduction to General Chemistry

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Introduction to General Chemistry

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  1. Introduction to General Chemistry Basic Principles

  2. Scientific Method • A systematic approach to research • Hypothesis: a tentative explanation for a set of observations and experiments • Law: a description of a phenomenon that allows for general predictions • Theory: a well-established explanation for scientific data; not fully tested; can be disproven • Experiments: systematic observations and measurements performed under controlled conditions

  3. What is Chemistry? • Chemistry is a scientific discipline that studies the inner properties of matter • Chemistry studies reactions between different substances and formation of new molecules


  5. What is matter? • Matter is any physically present substance that has a mass and occupies space • Matter is composed of atoms and molecules • Chemistry studies changes that occur in matter at the atomic and molecular level

  6. Classification of Matter

  7. Examples • Copper wire • A cup of tea • Oil / petroleum oil • Sea water • İron • Steel (alloy of iron and carbon)

  8. Units and Measurement • A measured quantity is given with a number and a unit • There are two common types of units used in science; metric units and SI units; SI units were derived from metric units but contain only one metric unit for each quantity

  9. Prefixes Used with SI Units 1 mL = 0.001 L 1 Mbytes = 1000000 bytes 1 nm = 0.000000001 m

  10. Common Measurements in Chemistry • Mass:measure of the amount of matter in a sample (kg, g) – mass is constant everywhere; (weight is not same as mass and it is the force that is applied on a object by gravity) • Volume: SI unit is m3, liter (L) is also used commonly; 1 L = 1 dm3 • Temperature:average kinetic energy of the atoms or molecules in a substance

  11. Common Measurements in Chemistry • Energy: capacity to perform a work, calorie (cal) and joule (J), 1 cal = 4.18 J • Density: mass per volume density = mass / volume “ g/mL “ is the unit commonly used

  12. Scientific Notation • In order to deal correctly with extremely large or extremely small numbers, we use scientific notation ( N × 10n, N is a number between 1 and 10) • Examples: 548.736 5.49 × 102 0.0000345 3.45 × 10-5

  13. Significant Figures • What is significant figures? • Meaningful digits in a measured quantity • Often impossible to report the exact value of a quantity from a measurement – there is always some error • The margin of error (uncertainty) should be indicated clearly – this is done by indicating the number of significant figures • The last digit is meant to be uncertain – indicates the error margin • Rules for determining significant figures: • 1) All non-zero digits are significant (both before and after decimal point) • 2) Zeros to the left of the first non-zero digit aren’t significant • 3) If the number ends with zeros at the right of the decimal point, those zeros are significant • 4) If a number ends in zeros to the left of the decimal point, those zeros might be significant or not (writing in exponential form solves this uncertainty) • When we report a set of numbers, there should be consistency in the number of significant figures

  14. Calculations with significant figures • Addition and subtraction: When adding/subtracting numbers having different number of significant figures, the result must be reported with the same number of decimal places as the original number that has the fewest decimal places (rounding-off) • Multiplication and division: The result of multiplication or division can’t have more significant figures than the original number with fewest significant figures • Exact numbers (e.g. number of objects, unit definitions) are considered to have infinite number of significant figures

  15. Accuracy and Precision • Accuracy: how close is a measurement to the true value of the measured quantity • Precision: how close are two or more measurements of the same quantity match each other

  16. Elements • An element is a singlesubstance in its simplest form that cannot be split into any more separate substances by chemical means • Everything around us is comprised of chemical elements • 112 elements, 90 of them naturally occurring • Only 2 elements are liquid at RT (bromine and mercury) and 11 are gases; all the rest are solids at RT

  17. Periodic Table • Periodic table is a chart in which elements with similar physical and chemical properties are grouped in a periodic way • The elements are arranged according to their atomic number • In a periodic table, horizontal rows are called periods and vertical groups are called groups • Elements within each group have similar chemical and physical properties • Groups 1-2 and 13-18 are called main group elements (also called 1A through 8A groups) • Groups 3-12 are called transition group elements (also called 3B through 12B group elements)

  18. Periodic Table

  19. Classification of Elements 17 nonmetals 8 semimetals

  20. Descriptive Names for Groups in the Periodic Table • 1A – alkali metals: lithium, sodium, potassium are most common, very reactive against air and water, hydrogen (H) is also in this group, but it is not a metal • 2A – alkaline earth metals: magnesium and calcium are the most abundant in nature among the group, found mainly as minerals • 7A – halogens: fluorine, chlorine, bromine, iodine are the most common – halogens react readily with metals to form salts (sodium chloride, calcium chloride) • 8A – noble gases: helium, neon, argon, krypton, xenon, radon – they are very unreactive gases – also called inert gases, they are present in monoatomic form • Transition metals (B group elements) – contain many of the common metals, such as iron, nickel, copper, cobalt, zinc, platinum, gold, silver

  21. Compounds • Compounds are pure substances that consist of two or more atoms of different elements held together by covalent or ionic bonds • The smallest structural unit of a compound is molecule • Compounds can be separated into smaller parts by chemical reactions • Compounds have a defined chemical structure (fixed ratio of atoms) – molecular formula • Examples;water (H2O) is a compound made of H2O molecules • Table salt (NaCl) is a compound made by NaCl units (ionic network)

  22. Changing of The Three States of the Matter • The extent of physical (non-covalent) interactions between molecules determines the physical state of a substance; solid, liquid, gas • Solid > liquid > gas (the order of the strength of non-covalent interactions) • The physical state of a substance can be changed by altering the number of non-covalent interactions between its molecules; this is achieved by giving or taking energy from the substance – generally by heat energy

  23. Three States of Matter • As we increase the energy of a substance, its molecules exhibit greater degree of movement and finally overcome the attractive forces holding the molecules together • Polar molecules have higher melting and boiling points, non-polar molecules have lower melting and boiling points (water b.p. =100 °C vs. methane b.p.= -161 °C)

  24. Atomic Theory • Theories of atom goes back to ancient Greek (Democritus) • First modern atomic theory is by John Dalton in 1808 • Summary of Dalton’s Atomic Theory: • 1) Elements are composed of very small particles, called atoms • 2) All atoms of an element are identical, but they are different than atoms of other elements • 3) Compounds are composed of atoms of more than one element. The ratio of numbers of different atoms in a compound is an integer or simple fraction (law of definite proportions) • 4) A chemical reaction includes separation, combination or rearrangement of atoms, not their creation or destruction

  25. Atoms and Atomic Theory • The smallest unit (particle) of an element is atom • Atom is made up of subatomic particles; protons, neutrons and electrons – number of these determine the characteristic of an atom

  26. Atoms • Atomic Number (Z):number of protons (or electrons in a neutral atom) in an atom • Mass Number (A): (number of protons) + (number of neutrons)


  28. Atoms • Ions: gain electron – anion, lose electron – cation(fluoride, F- ; sodium cation, Na+) • Atoms form ions as part of their reaction with other atoms to form molecules. • It is the movement and sharing of electrons between atoms and molecules that lies at the heart of virtually all chemical processes–including those occurring in the cells of your body. • So the readiness with which an atom gains or loses electrons dictates its reactivity

  29. Atoms • Isotopes:atoms with the same number of protons and electrons, but different number of neutrons 16O, 17O, and 18O • Elements are present in nature as mixtures of their isotopes