Chemistry

1. Chemistry

2. Matter • Organisms are composed of matter • Matter is anything that takes up space and has mass • Matter is composed of chemical elements • Matter is found on the Earth in three physical states • Solid • Liquid • Gas

3. States of Matter • Gases take the shape and volume of their container and can be compressed to form liquids. • Liquids take the shape of their container, but they do have their own volume • Solids are rigid and have a definite shape and volume.

4. Sodium Chlorine Sodium chloride Classification of Matter • Element: a substance composed of only one type of atom (all the atoms have the same number of protons). • Molecule: a unit composed of two or more atoms joined together by chemical bonds • Compound: a substance composed of 2 or more elements that have been joined by chemical bonds • Mixture: a combination of 2 or more substances that do NOT chemically bond e.g. sugar mixed with salt

5. Pure Substances and Mixtures

6. Elements • If a pure substance cannot be decomposed into something else, then the substance is an element • There are 114 elements known, 92 naturally occurring • Each element is given a unique chemical symbol (one or two letters) Periodic Table

7. Essential Elements Of Life • Only about 25 of the elements are essential to life • Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter • Most of the remaining 4% consists of calcium, phosphorus, potassium, and sulfur • Trace elements are those required by an organism in minute quantities

8. Essential Elements Of Life

9. Periodic Chart

10. Atoms • Each element consists of one kind of unique atom • An atom is the smallest unit of matter that still retains the properties of an element, it cannot be broken down to other substances by chemical reactions

11. Subatomic Particles Atoms are composed of subatomic particles Relevant subatomic particles include: Neutrons (no electrical charge) Protons (positive charge) Electrons (negative charge) Neutrons and protons form the atomic nucleus Electrons form a cloud around the nucleus Nucleus (a) (b) Cloud of negative charge (2 electrons) 2 Protons Neutrons 2 Electrons 2

12. Atomic Number And Atomic Mass • Atoms of the various elements differ in number of subatomic particles • An element’s atomic number is the number of protons • The number of protons (atomic number) determines the element’s properties • An element’s mass number is the sum of protons plus neutrons in the nucleus • Atomic mass, the atom’s total mass, can be approximated by the mass number

13. Atomic number Element symbol Mass number Periodic Chart

14. Orbitals • Electrons orbit the nucleus of an atom in specific electron shells • Each Orbital holds a maximum of 2 electrons each • Several orbitals may be the same distance from the nucleus and thus contain electrons of the same energy. Such electrons are said to occupy the same energy level or shell. • Rule of Eights for filling each shell: First electron shell (can hold 2 electrons) Outermost electron shell (can hold 8 electrons) Electron Hydrogen (H) Atomic number = 1 Carbon (C) Atomic number = 6 Nitrogen (N) Atomic number = 7 Oxygen (O) Atomic number = 8

15. Electron Shell Significance • Electrons determine how an atom behaves when it encounters other atoms • Outer orbital (valence shell) determines reactivity of atom - Electronegativity • Atoms “desire” full outer orbitals • Give up electrons (Na) • Take electrons (Cl) • Share electrons (O2) • Noble gases - full outer shells (inert)

16. Chemical Bonding and Molecules • Chemical reactions enable atoms to give up or acquire electrons in order to complete their outer shells • These interactions usually result in atoms staying close together • The atoms are held together by chemical bonds

17. Kinetic Theory Of Matter • All atoms and molecules are in constant random motion. (Energy of motion is called kinetic energy.) • The higher the temperature, the faster the atoms and molecules move. • All motion theoretically stops at absolute zero.

18. Energy • Energy is the capacity to do work or ability to cause change. Any change in the universe requires energy. Energy comes in 2 forms: • Potential energy is stored energy. No change is currently taking place • Kinetic energy is currently causing change. This always involves some type of motion.

19. On the platform, a diver has more potential energy. Diving converts potential energy to kinetic energy. Climbing up converts kinetic energy of muscle movement to potential energy. In the water, a diver has less potential energy. Forms Of Energy • Kinetic energy is the energy associated with motion • Potential energy • Is stored in the location of matter • Includes chemical energy stored in molecular structure • Energy can be converted from one form to another • First Law Of Thermodynamics states that energy cannot be created or destroyed; energy can be transferred or transformed

20. Temperature, Pressure, And Volume • Volume – Pressure Relationship • At a constant temperature, volume is inversely proportional to pressure • Volume – Temperature Relationship • At constant pressure, the volume is directly proportional to temperature

21. Chemical Reactions • Cells constantly rearrange molecules by breaking existing chemical bonds and forming new ones • Such changes in the chemical composition of matter are called chemical reactions • Chemical reactions enable atoms to give up or acquire electrons in order to complete their outer shells • These interactions usually result in atoms staying close together • The atoms are held together by chemical bonds • Reactions can be written as equations

22. Hydrogen gas Oxygen gas Water Products Reactants Chemical Equations • The chemical equation for the formation of water can be visualized as two hydrogen molecules reacting with one oxygen molecule to form two water molecules: • 2H2 + O2  2H2O Reactants

23. Reading Chemical Equations • The plus sign (+) means “react” and the arrow points towards the substance produce in the reaction • The chemical formulas on the right side of the equation are called reactants and after the arrow are called product • The numbers in front of the molecules or atoms indicate the number of individual molecules or atoms (stoichiometric coefficients) • The numbers behind are subscripts indicating the molecules or atoms are bonded 2Na + 2H2O  2NaOH + H2 Reactants Products

24. Chemical Reactions • Are dependent on : • Concentration • Speed • Energy (energy of activation) • Orientation

25. Types Of Chemical Reactions • Synthesis reactions - atoms or molecules combine to form a product • Decomposition reactions - molecules breakdown into smaller molecules or atoms • Exchange reactions - molecules exchange constituent components (swap partners) • Reversible reactions - the product of a previous reaction can revert to the original reactants.

26. Combination (Synthesis) Reactions • Combination (Synthesis) reaction occurs when two or more substances react to form products: Na + Cl  NaCl Ca + 2NaCl  CaCl2 • In both cases, Sodium and Calcium combine with Chlorine

27. Decomposition Reactions • Decomposition reaction is when one substance undergoes a reaction to produce two or more substances: 2H2O  2H2 + O2 H2O2 H2O + O2

28. Exchange Reactions • Exchange reaction occurs when molecules “swap partners”: NaOH + HCl  NaCl + H2O H2CO3 + NaOH  H2O + NaHCO3

29. Reversible Reactions • Reversible reactions can go forwards (decomposition) or backwards (combination): H2CO3  CO2 + H2O • Chemical Equilibrium is defined as the state of dynamic balance in which the rates of forward and reverse processes (reactions) are equal

30. Chemical Products • Element: a substance composed of only one type of atom (all the atoms have the same number of protons). • Molecule: a unit composed of two or more atoms joined together by chemical bonds • Compound: a substance composed of 2 or more elements that have been joined by chemical bonds • Mixture: a combination of 2 or more substances that do NOT chemically bond e.g. sugar mixed with salt

31. Ionic Bonds • Atoms sometimes strip electrons from their bonding partners • An example is the transfer of an electron from sodium to chlorine • After the transfer of an electron, both atoms have charges • A charged atom (or molecule) is called an ion • An anion is a negatively charged ion • A cation is a positively charged ion • An ionic bond is an attraction between an anion and a cation - oppositely charged ions

32. Ions And Ionic Compounds • When an atom or molecule loses electrons, it becomes positively charged. • For example, when Na loses an electron it becomes Na+. • Positively charged ions are called cations. • When an atom or molecule gains electrons, it becomes negatively charged. • For example when Cl gains an electron it becomes Cl-. • Negatively charged ions are called anions. • An atom or molecule can lose more than one electron. • When molecules loose electrons, polyatomic ions are formed.

33. Na+ Cl– Ionic Compounds • Compounds formed by ionic bonds are called ionic compounds, or salts • Salts, such as sodium chloride (table salt), are often found in nature as crystals

34. Hydrogen atoms (2 H) Hydrogen molecule (H2) Covalent Bonds • Molecules are formed by covalent bonds • A covalent bond is when two atoms share one or more pairs of outer-shell electrons (valence electrons) • In a covalent bond, the shared electrons count as part of each atom’s valence shell • Much stronger than ionic bonds – holds lots of Energy • A single covalent bond, or single bond, is the sharing of one pair of valence electrons • A double covalent bond, or double bond, is the sharing of two pairs of valence electrons • Covalent bonds can form between atoms of the same element or atoms of different elements

35. Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Oxygen (O2) Covalent Bonds

36. Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Water (H2O) Methane (CH4) Covalent Bonds

37. Covalent Bonds Figure 2.9

38. – O H H + + H2O Electronegativity • Outer orbital (valence shell) determines reactivity of atom - Electronegativity • Electronegativity is an atom’s attraction for the electrons in a covalent bond • The more electronegative an atom, the more strongly it pulls shared electrons toward itself

39. Polar Covalent Bond • In a nonpolar covalent bond, the atoms share the electron equally • In a polar covalent bond, one atom is more electronegative, and the atoms do not share the electron equally

40. () () () () The Structure Of Water • Its two hydrogen atoms are joined to one oxygen atom by single covalent bonds • But the electrons of the covalent bonds are not shared equally between oxygen and hydrogen • This unequal sharing makes water a polar molecule Unnumbered Figure 2.2

41. () Hydrogen bond () () () () () () () (b) Figure 2.11b Hydrogen Bonds • A hydrogen bond forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom • In living cells, the electronegative partners are usually oxygen or nitrogen atoms

42. – + Water (H2O) + Hydrogen bond – Ammonia (NH3) + + + Hydrogen Bonds

43. Weak Chemical Bonds • Most of the strongest bonds in organisms are covalent bonds that form a cell’s molecules • Weak chemical bonds, such as ionic bonds and hydrogen bonds, are also important • Weak chemical bonds reinforce shapes of large molecules and help molecules adhere to each other

44. Biological Importance Of Water • Acts as a powerful solvent • Participates in chemical reactions • Water has a high specific heat which moderates temperature - absorbs and releases heat very slowly, minimizes temperature fluctuations to within limits that permit life • Heat is absorbed when hydrogen bonds break • Heat is released when hydrogen bonds form • Requires a great amount of heat to change to a gas • Heat of vaporization - the quantity of heat a liquid must absorb for 1 gram of it to be converted from a liquid to a gas • Evaporative cooling - Allows water to cool a surface due to water’s high heat of vaporization • Acts as a lubricant

45. Polarity & Hydrogen Bonds • Cohesion - molecules attract other water molecules • Capillarity • Water molecules are drawn up a narrow tube • Helps pull water up through the microscopic vessels of plants • Surface tension • water molecules on the surface cling to each other – related to cohesion • Is a measure of how hard it is to break the surface of a liquid • Adhesion - water molecules attract other charged substances

46. Negative oxygen regions of polar water molecules are attracted to sodium cations (Na+). Na+ – + + – + Na+ – – Positive hydrogen regions of water molecules cling to chloride anions (Cl–). – + + Cl – Cl– + – – + – + – – Water As A Solvent • Water is a versatile solvent due to its polarity • It can form aqueous solutions • The different regions of the polar water molecule can interact with ionic compounds called solutes and dissolve them

47. Formula and Molecular Weights • Formula weights (FW) is the sum of the atomic weights of each atom in the chemical formula. • FW (H2SO4) = 2AW(H) + AW(S) + 4AW(O) • = 2(1.0 amu) + (32.0 amu) + 4(16.0) • = 98.0 amu • If the chemical formula is also its molecular formula then the weight is called the molecular weight (MW). • MW(C6H12O6) = 6(12.0 amu) + 12(1.0 amu) + 6(16.0 amu) =????

48. The Mole • The unit we use to express the quantity of atoms, ions, and molecules that an object contains is called mole. • Mole: convenient measure chemical quantities. • The actual number of atoms, ions, or molecules in 1 mole of something = 6.0221367  1023 (Advogadro’s number). • Thus, • 1 mole of 12C atoms = 6.02 x 1023 12C atoms • 1 mole of H2O molecules = 6.02 x 1023 molecules • 1 mole of NO3- ions = 6.02 x 1023 ions

49. Visualizing The Mole Concept Different Units

50. Solution Composition • Solutions are homogenous mixtures of two or more substances: • Solute: present in smallest amount and is the substance dissolved in the solvent. • Solvent: present in the greater quantities and is used to dissolve the solute. • Example: NaCl dissolved in Water (water = Solvent and NaCl = Solute) • Change concentration by using different amounts of solute and solvent. • Molarity: Moles of solute per liter of solution.