Biology 1000 Lecture #2-2 (1) PDF

Lecture #2 Basic Chemistry  Chemical Composition of Living Organisms All living things are composed of matter Matter always occupies space and has a mass Examples: rock, water, grass, animals, human beings All matter is composed of elements Ele...

Lecture #2 Basic Chemistry  Chemical Composition of Living Organisms All living things are composed of matter Matter always occupies space and has a mass Examples: rock, water, grass, animals, human beings All matter is composed of elements Elements cannot be further broken down chemically There are 92 elements that are known to be found in nature Examples: carbon, nitrogen, oxygen, hydrogen, gold, copper Each element is given a symbol Examples: Carbon (C), Oxygen (O), Sodium (Na)  Chemical Composition of Living Organisms There are 25 elements necessary for human life 96% of the bodies weight can be attributed to carbon, hydrogen, oxygen and nitrogen These four elements are the main ingredients of sugars, fats, proteins and nucleic acids The remaining 4% of body weight is comprised of: Calcium Magnesium Chlorine Sodium Potassium Phosphorous Sulfur Trace elements are essential for life but are required in very small amounts  Trace Elements Iron is a trace element that accounts for only 0.004% of body mass It is absolutely necessary because it is a key component of hemoglobin which is used to transport oxygen in blood Iodine is another trace element that is required for human life: 0.15mg of iodine must be acquired each day Iodine is a key component of hormones produced by the thyroid gland Iodine deficiency leads to an enlargement of the thyroid gland called a goiter To avoid this condition, table salt in many countries is supplemented with iodine Referred to as iodized  Compounds Elements combine to form compounds (molecules): A compound is a substance consisting of two or more elements in fixed proportion Much more common than free elements Examples: NaCl (sodium chloride), H2O (water) Compounds have entirely different properties than the constituent elements Example: Sodium (Na) is a metal, Chlorine (Cl) is a gas, combined they produce table salt which is edible  Compounds Most compounds found in living organisms contain carbon (C), oxygen (O), Nitrogen (N) and Hydrogen (H): Sugars are composed of carbon, oxygen and hydrogen Proteins are composed of carbon, oxygen, hydrogen and nitrogen as well as a small amount of Sulfur Different arrangements and different proportions of these elements provide proteins and sugars with characteristic properties which are very different from one another  Atoms Atoms are the smallest units of matter that still display properties of the element Example: a carbon atom is the smallest unit that still has properties of carbon Atoms consist of subatomic particles: Protons: carry a positive charge and are found in the nucleus Neutrons: do not carry any charge and are found in the nucleus Electrons: carry a negative charge and are found outside of the nucleus in orbitals The nucleus is the central core of the atom The electrons are kept near to the nucleus because their negative charge is attracted to the positively charged core (nucleus)  Atoms The main difference between different atoms is the number of sub- atomic particles All atoms of a particular element have the same number of protons within their nucleus The atomic number of the element is the number of protons within the nucleus of each atom Example: Carbon has an atomic number of 6 therefore each carbon atom has 6 protons within the nucleus Each atom has an equal number of protons and electrons when it is neutral therefore the atomic number is also equal to the number of electrons The mass number of an atom is equal to the sum of the protons and the neutrons found within the nucleus  Atoms Isotopes: All atoms of an element always have the same atomic number however the mass number may differ The term isotope refers to an atom that has the exact same number of protons and electrons as all other atoms of an element but differs from other atoms in the number of neutrons present within the nucleus Since the number of protons and electrons remain unchanged, the isotope has identical chemical properties to the other atoms of the element There are three naturally occurring carbon isotopes Carbon-12 accounts for 99% of all carbon found in nature  Electron Arrangement of an Atom Electrons determine the chemical reactivity of an atom Different electrons have varying degrees of energy The further an electron is from the nucleus the greater the energy it has Electrons occur only at certain energy levels referred to as electron shells The number of electron shells surrounding the nucleus varies from atom to atom depending on the number of electrons  Electron Arrangement of an Atom With some exception outer electron shells can hold 8 electrons If the outer shell has fewer than 8 electrons the element is considered reactive The fewer electrons needed to achieve a number of 8 the more reactive the atoms of the element are To achieve a number of 8 electrons in the outer shell atoms form bonds with other atoms either via: Electron sharing: covalent bonds Electron Donation: ionic bonds  Ionic Bonds An atom that has only one electron in its outer shell will donate the electron rather than looking to gain or share 7 The electron will be donated to an atom that is looking to acquire an electron because it has 7 electrons in its outer shell Example: NaCl Sodium has one electron in its outer shell and chlorine has 7 electrons in its outer shell Sodium donates the electron to chlorine so that sodium has 0 electrons in its outer shell and chlorine has 8 electrons in its outer shell This results in the formation of ions (‘ionic’ bonds): Sodium lost an electron and was previously neutral (now Na+) Chlorine gained an electron and was previously neutral (now Cl-) The bond formed is a result of the attraction between the newly formed positive and negative charges  Covalent Bonds These bonds form when two atoms share electrons Molecules are formed when two or more atoms share electrons Example: hydrogen (H2) is joined by a covalent bond When two electrons are shared, a single bond is formed Example H2 H-H When 4 electrons are shared a double bond is formed Example: O2 O=O The number of covalent bonds that an atom can form is equal to the number of electrons needed to fill its outer shell Example: Carbon requires four electrons to achieve eight electrons in its outer orbital therefore it can form four covalent bonds  Unequal Electron Sharing Electrons shared in covalent bonds are involved in a ‘tug-of-war’ The atoms sharing the electrons are both pulling on the electrons involved in the bond Electronegativity is a measure of an atoms attraction for electrons The greater the electronegativity of an atom the more it will pull electrons it’s way  Unequal Electron Sharing Molecules of only one atom such as O2 will equally share electrons Covalent bonds where electrons are shared equally are called non-polar covalent bonds These bonds form in between atoms of similar electronegativity Non-polar covalent bonds can also form in between different atoms that are of similar electronegativity Example: methane (CH4)  Unequal Electron Sharing Covalent bonds that form between atoms of very different electronegativities result in unequal electron sharing These bonds are called polar covalent bonds The more electronegative atom pulls the electrons its way more than the less electronegative atom Unequal electron sharing causes a partial positive charge on the less electronegative atom and a partial negative charge on the more electronegative atom Oxygen is one of the most electronegative elements Example: H2O  Hydrogen Bonds Bonds that form between molecules and within molecules are important for cell function These bonds are non-covalent and weaker than covalent bonds Hydrogen involved in a polar covalent bond will carry a partial positive charge As a consequence the hydrogen will orient itself near to an atom in an adjacent molecule that has a partial negative charge This weak non-covalent bond is called a hydrogen bond because the positive molecule involved is always hydrogen Hydrogen Bonds and Temperature Heat is defined as the amount of energy associated with the movement of atoms and molecules in an object Temperature measures heat intensity: the measurement of the average speed of molecules Heat is required to break hydrogen bonds formed between water molecules Heat is released when these hydrogen bonds form Since temperature measures the speed of molecules, heat must be added to water in order to break hydrogen bonds Once these bonds are broken the molecules can move much faster and as a result the temperature of the water will increase Hydrogen Bonds and Temperature When water is cooled heat is released as hydrogen bonds are reformed between water molecules This results in a decreased speed of the molecules leading to a decrease in temperature  Evaporation Evaporation of a substance moderates temperature The molecules with the greatest energy (the hottest) leave the substance The remaining liquid is cooler as a result of this loss Example: sweating Boiling water  Hydrogen Bonds and Density Water exists in three primary states: Solid Liquid Gas (water vapor) Solids are always more dense than gas where density is the number of particles (molecules) per unit area Exception: water Ice has a smaller density than water vapor Explains why ice cubes float in water Due to hydrogen bonds When water freezes each water molecule forms four bonds with adjacent water molecules result in the creation of a crystal These hydrogen bonds are extremely stable In liquid water the hydrogen bonds are less stable and continually break and reform  Water as a Solvent A solution is a liquid consisting of a uniform mixture of two or more substances The solution consists of a: Dissolving agent called a solvent: usually water A dissolved substance called a solute: examples include sugar and salt When water is the solvent the solution is referred to as an aqueous solution Water is such a great solvent because of the polarity of the molecules NaCl in water dissociates forming Na+ and Cl- ions These ions are then surrounded by water molecules The partial positive charge on water’s hydrogen interacts with the Cl- ions One of the partial negative charges on water’s oxygen interacts with the Na+ ions  pH Scale Measured on a scale of 0-14 A solution with a pH of 7 is considered neutral The amount of H+ in solution is equal to the amount of OH- in solution Solutions with a pH of 0-6 are acidic pH= 0 is the most acidic pH= 6 is the least acidic (near neutral) Solutions with a pH of 8-14 are basic pH= 8 is the least basic pH= 14 is the most basic  Acids and Bases Acids: Dissociate when placed in water Release H+ ions Example: HCl H + + Cl – The more HCl that is added to water, the more H + will be present following dissociation The greater the amount (concentration) of H + in the solution, the lower the pH of the solution Lower pH means a greater acidity Bases: Dissociate in water also Release OH – ions Example: NaOH Na + + OH – The OH - that is generated from the dissociation of NaOH binds to H + The greater the amount (concentration) of OH - in the solution, the greater the pH of the solution Higher pH means that the solution is more basic