The chemical nature of cells

1. The chemical nature of cells (all that tricky chemistry stuff!!!)

2. Introduction • At the most basic level, the cell is a highly organised assemblage of atoms interacting with each other in a myriad of chemical reactions. • In order to understand the molecular nature of the cell, we need to revisit some chemical concepts.

3. Some definitions • Matter • Anything that takes up space and has mass. • Atom • An atom a fundamental piece of matter. Everything in the universe (except energy) is made of matter, and, so, everything in the universe is made of atoms. • An atom itself is made up of three tiny kinds of particles called subatomic particles: protons, neutrons, and electrons. • Isotope • Different forms of atoms of the same element. They have the same number of protons in their nuclei but a different number of neutrons.

4. Some more definitions • Molecule • Formed when two or more atoms are joined chemically. • Element • An element is a substance consisting of only one type of atom. e.g oxygen gas is O2 (pairs of oxygen atoms combine to give oxygen gas) • Compound • A compound is a molecule that contains at least two different types of molecules. e.g. water is H2O (two hydrogens and one oxygen)

5. More about elements • There are 92 naturally occurring elements. • Only 11 of these are found in organisms in more than trace amounts, and four of these make up more than 99% of organisms by weight. • The four most common elements in living organisms are carbon (C), hydrogen (H), oxygen (O) and nitrogen (N).

6. Memory Aid!!! Most common elements: CHONPS • C for carbon • H for hydrogen • O for oxygen • N for nitrogen • P for phosphorous • S for sulfur

7. More about atoms • The red and green circles in the centre are the proton and neutron, they both make up the nucleus. • The blue circle is the electron and the black ring shows its orbit around the nucleus.

8. Even more about atoms – particularly electrons and electron shells • The orbit or region of space around the nucleus in which electrons are found is referred to as an electron shell. • There are rules as to the number of electrons that can be held in each shell. • The arrangement of electrons within these shells is called the electron configuration. • The electrons that occupy the outermost shell are called valence electrons. • The chemical behaviour of an atom is a function of its arrangement of electrons – in particular, the number of valence electrons in its outermost electron shell. • Atoms are most stable when their outermost shell is full.

9. Making molecules • Atoms are most stable when their outermost electron shell is full. • In order to achieve this state, atoms tend to combine with other atoms to form molecules. • There are two types of chemical bonds that can hold the atoms within a molecule together. These are ionic bonds or covalent bonds.

10. Ionic Bonds • Atoms gain or loseelectrons in order to increase stability. • Atoms that lose electrons have a positive charge and are called cations. • Atoms that gain electrons have a negative charge and are called anions. • An ionic bond is an electrical attraction between two oppositely charged atoms or groups of atoms.

11. Example of ionic bonding NaCl or sodium chloride • Sodium has 1 valence electron which it loses to become a cation with a charge of 1+ • Chlorine has 7 valence electrons and acquires 1 extra electron from sodium in order to become an anion with a charge of 1- • The electrostatic attraction between positive sodium and negative chloride (name changes when it forms an ion) holds the molecule together.

12. Covalent Bonds • Atoms can share electrons in order to increase their stability • Covalent bonding is where the atoms share pairsof outer shell or valence electrons. • Covalent bonding may be single or multiple, depending on the number of pairs the atoms share. • Sometimes in covalent bonds, one atom attracts the shared electrons more strongly than the other, resulting in a polar covalentbond.

13. Example of covalent bonding H2O or water • Hydrogen (H) has 1 valence electron but needs a total of 2 in order to be a stable atom. • Oxygen has 6 valence electrons but needs a total of 8 in order to be a stable atom. • By sharing electrons, each hydrogen atom has two valence electrons, thus filling their outer orbits. Likewise, oxygen now has 8 outer orbit electrons. • This makes for a good chemical bond and a stable molecule.

14. Polar Covalent Bonds • Polar covalent bonds are formed when one atom attracts the shared electrons more strongly than the other. • Molecules that contain polar covalent bonds are referred to as polar molecules and molecules that are formed by covalent bonds but don’t have polar bonds are called non-polar molecules. • The measure of the ability of an atom to attract electrons in a bond is called electronegativity. • The region of the molecule which contains the atom with the greatest electronegativity will have a slightly negative charge compared to the rest of the molecule. The rest of the molecule will have a slightly positive charge. This opposite charge separated by a distance is called a dipole.

15. Electronegativity • The higher the value on the Pauling Electronegativity Scale, the stronger the atom’s electron attracting power.

16. Simple analogy for electronegativity • Sharing electrons in covalent bonds is like trying to get a couple to share the doona equally. • Someone is always going to have more pulling power! • One atom is always going to “hog” the electrons. • In a water molecule (H2O), the oxygen atom will always attract the shared electrons more than the hydrogen atoms do.

17. Interactions between molecules • Intermolecular bonds are important in maintaining the 3D structure of large biomolecules such as proteins and nucleic acids. • These interactions also allow a molecule to bind specifically but transiently with another molecule. • Individual interactions are weak and constantly break and reform at the physiological temperatures of organisms. • Multiple interactions act together to produce highly stable and specific associations between parts of a large biomolecule and/or between different molecules. • Four types of interactions are responsible for intermolecular bonds: • Hydrogen bonds • Van der Waals interactions (transient dipoles) • Hydrophobic bonds • Ionic interactions

18. Hydrogen bonds(dipole-dipole bonding) • Weak electrostatic attraction between the negative region (δ-) of one polar molecule and the positive region (δ+) of another polar molecule.

19. Van der Waals interactions (transient dipoles) • At very short distances all atoms and molecules show a weak bonding interaction due to their fluctuating electrical charges. • Electrical charges fluctuate due to the uneven distribution of electrons as they orbit the nucleus of an atom. The larger the atoms involved the greater the fluctuation in charge.

20. Hydrophobic bonds • Non-polar molecules such as fats or oils will aggregate together when placed in a polar substance such as water. This aggregation is referred to as a hydrophobic bonds. • It is not a separate bonding force but is due to water molecules excluding the non-polar molecules, forcing them to adhere to one another. Red molecules are H2O

21. Ionic interactions • Ionic compounds are generally soluble in water due to ionic interactions with water molecules. • Atoms in functional groups can donate or accept protons (H+) forming ions that interact with other charged groups on atoms on different molecules. Molecule with NH2 group attached has gained H+ so becomes positive ion. Molecule with COOH group attached has lost H+ so becomes negative ion.

22. Inorganic and Organic Molecules • Both living and non-living things are made from the same chemical elements but their is a difference in the way that these elements are put together to make larger molecules. • Organic compounds contain carbon and hydrogen (and sometimes other elements such as oxygen and nitrogen). They are called organic compounds as the first ones studied were produced by or found in living organisms. • All other compounds are called inorganic compounds.

23. Biologically ImportantInorganic Molecules • Inorganic molecules important for living organisms include: • nitrogen – present in all proteins and nucleic acids. Fixed from the atmosphere by nitrogen-fixing bacteria • minerals – found in the cytosol and structural components of cells and in the molecules of enzymes and vitamins. Important minerals include phosphorous, potassium, calcium, magnesium, iron, sodium, iodine and sulfur.Examples: phosphorous present in phospholipids of cell membrane and in ATP, magnesium important component of chlorophyll, iron an important component of haemoglobin • oxygen – most cells require oxygen to release usable energy from food molecules • carbon dioxide – main source of carbon for the production of organic molecules despite the fact that it only makes up 0.035% of atmosphere by volume. Organic molecules (sugars) are eaten by animals and carbon dioxide released back into atmosphere as an end-product of cellular respiration. • water - chemical reactions in cells occur in a water environment

24. More about carbon • All the chemicals of life on this planet, with the exception of water, are based on the carbon atom. • It is the valency of carbon that allows it to form the base of all chemicals of life. The carbon atom has four valence electrons, meaning it can form four stable covalent bonds with other atoms. • Carbon atoms can also bond with other carbon atoms to form straight and branched chain and ring structures of various sizes and complexity that form the backbone to many biological molecules. • Carbon atoms can share more than one pair of electrons between two carbon atoms, resulting in the formation of double and triple bonds. • Molecules containing only carbon and hydrogen atoms are known as hydrocarbons. Hydrocarbons are non-polar and hence insoluble in water. • What creates the diversity and chemical properties of carbon based molecules is the addition of other groups of atoms to the hydrocarbon backbone. • Groups of atoms that confer water solubility and chemical properties to the hydrocarbon chain are known as functional groups. • These groups of atoms are more reactive than the hydrocarbon portion of biomolecules and often contain highly electronegative atoms which convey a polarity to their end of the molecule while the hydrocarbon chain is non-polar.

25. Water • Water covers about 75% of our planet’s surface and makes up 70% to 90% of the cell content of living things. • It has a number of unique properties that support life on Earth. • These properties are a direct result of the polar nature of the water molecule

26. Why is water so special?

27. Water as a pH buffer • pH is a measure of the hydrogen ions in a solution and hence the state of acidity or alkalinity of a solution. • pH scale is 0 to 14. • Less than 7 indicates higher concentration of H+ (acidic solution) • Greater than 7 indicates higher concentration of OH- (alkaline solution) • Pure water has pH of 7.0 and is a neutral solution but water readily ionizes or break ups to form H+ and OH- ions. • This allows the cellular fluids to balance changes in pH as binding of these ions to other substances within the cell prevents severe changes in the pH of a cell or fluid.

28. I’m a little water molecule!!!(to the tune of “I’m a little teapot”) I’m a little water molecule They call me H2O I’m not to good at sharing My H more positive than my O That’s why they call me polar Cause my molecules attract each other In between them, H bonds I’m a BIG SELF LOVER!