Level 6 Food Science: Chemistry Fundamentals

• Fundamentals of Science is a Level 6 Certificate in Food Science and Technology.

Module Content

• Biology topics include Cell Structure and Function, Proteins and Analysis, Lipids, Carbohydrates, Nutrition, Enzymes, Quantitative Analysis, Weights, Measures, and Preparation.
• Chemistry topics include Periodic Table, Chemical Bonds, Acids and Bases, and Elemental Analysis.

Session Objectives

• Identify main groups of elements in the Periodic Table.
• Calculate the number of electrons, protons, and neutrons for elements.
• Calculate the number of electrons in the valence shell.
• Predict the type of bonding.
• Draw Lewis structures for main elements.

Goal of Chemistry

• The goal is to understand how elements combine to form new materials
• The goal seeks to find the rules which apply to creating new materials, and when combination is impossible.

Chemistry Building Blocks

• Elements can combine to make molecules such as paracetamol.
• Small molecules can be pharmaceutical medicines, examples include aspirin and ibuprofen.
• Molecules can combine to create bigger molecules which include bio-molecules and polymers.
• Examples of bio-molecules/polymers: DNA, RNA, Carbohydrates, and Proteins.
• These examples are made from the same building block elements.

Elements

• Substances that can't be broken down into simpler substances.
• Molecules are formed by a combination of elements and serve as the foundation for larger structures.
• Small Molecules are specific molecules that serve as pharmaceutical compounds.
• Biomolecules are larger molecules essential for life, like DNA and proteins.
• Polymers are complex structures formed by repeating units of smaller molecules.

Periodic Table

• Chemistry relies on the periodic table of elements to understand how different elements interact and form compounds

Symbols of Elements

• Element symbols use 1 or 2 letter abbreviations.

• The first letter is always capitalized.

  • C represents carbon.
  • Co represents cobalt.
  • N represents nitrogen.
  • Ca represents calcium.
  • F represents fluorine.
  • Br represents bromine.
  • O represents oxygen.
  • Mg represents magnesium.

• Copper is Cu and its Latin name is Cuprum.

• Gold is Au and its Latin name is Aurum.

• Lead is Pb and its Latin name is Plumbum.

• Mercury is Hg and its Latin name is Hydrargyrum.

• Potassium is K and its Latin name is Kalium.

• Silver is Ag and its Latin name is Argentum.

• Sodium is Na and its Latin name is Natrium.

Body Composition

• Elemental composition of the human body is as follows:
  • Oxygen: 65%
  • Carbon: 18.5%
  • Hydrogen: 9.5%
  • Nitrogen: 3.2%
  • Calcium: 1.5%
  • Phosphorus: 1.0%
  • Potassium: 0.4%
  • Sulfur: 0.3%
  • Sodium: 0.2%
  • Chlorine: 0.2%
  • Magnesium: 0.1%
  • Iron: 0.006%

Atoms and Molecules

• An atom is the smallest part of an element.
• A molecule is the general term used to describe any atoms connected by chemical bonds.
• Every combination of atoms is a molecule.

Compounds and Substances

• A compound is a molecule made of atoms from two or more different elements that are bonded together.
• A substance has a definite fixed composition, and can be pure or a mixture.
• A substance can be an element or a compound.
• All compounds are substances, but not all substances are compounds.
• Water is a compound.
• Gold is a substance.

Subatomic Particles in Atoms

• Atoms are composed of a positively charged core called the nucleus.

• The nucleus contains protons and neutrons.

• A negatively charged exterior surrounds the nucleus, that is made up of electrons.

• Different elements are distinguished by different numbers of protons.

• Electron:

  • Symbol is e-
  • Charge is 1-
  • Mass is 0

• Proton:

  • Symbol is p+
  • Charge is +
  • Mass is 1

• Neutron:

  • Symbol is n
  • Charge is 0
  • Mass is 1

Compounds vs. Molecules

• A Molecule is formed when two or more atoms join together chemically (ie. H2)
• A Compound is a molecule that contains at least two different elements (ie. NaCl)
• All compounds are molecules, but not all molecules are compounds.

Organization in the Periodic Table

• The periodic table has vertical columns
• Elements in a column have similar chemical and physical properties.
• It is divided into metals and non-metals.
• Metal columns are divided into alkali, alkali earth, transition and other metals.

Metals vs Non Metals

• Blue represents the metals
• Red represents non-metals
• Green represents metalloid

Periodic Table Organisation

• Alkali metals are in the first column
• Alkali earth metals are in the second column
• Transition metals are in the middle section
• Other Metals are on the right of transition metals

Chemical Properties

• How does the element combine (react) with other elements.
• Elements combine by bonding.
• Bonds are formed by transferring or sharing electrons.
• The number of electrons determines the element loose or gain.
• Whether the element is toxic or have flammability.
• Is the element acidic.

Physical Properties

• Physical properties consist of:
  • Colour
  • Size
  • Shape
  • Density
  • Freezing and Boiling Points
  • Smell

Metals vs. Nonmetals

• Metals combine with other elements by losing electrons.
• They are located to the left of the periodic table.
• Metals are shiny and ductile.
• Good conductors of heat and electricity.
• Nonmetals combine with other elements by gaining or sharing electrons.
• They are located to the right in the periodic table.
• Nonmetals are dull and brittle.
• They are poor conductors and good insulators.

Chemical Properties: Atomic Number

• Depends on the atomic number.
• This is the number of protons.
• Mass number is the number of protons and neutrons.
• Charge on the atoms depends on the number of electrons which can change in a chemical reaction.
• Atoms can either gain or lose electrons.

Subatomic Particles Examples

• Oxygen
  • 16O
  • 8
  • 8 p+
  • 8 n
  • 8 e-
• Phosphorus
  • 31P
  • 15
  • 15 p+
  • 16 n
  • 15 e-
• Zinc
  • 65Zn
  • 30
  • 30 p+
  • 35 n
  • 30 e-

Number of Electrons and Charge

• An atom is neural when the number of protons and electrons are equal, meaning that the net charge is zero.
• Number of protons = Number of electrons.
• Atomic number = Number of electrons
• If the atom gains an electron it becomes negative.
• If the atom loses an electron it becomes positively charged.

Isotopes

• Isotopes are atoms with same number of protons but different numbers of neutrons
• Isotopes are atoms of the same element (same atomic number) with different mass numbers.
• A stable isotope is not radioactive.
• Number of electrons is the same
• Number of protons is the same
• Number of neutrons is different
• Mass is different

Isotopes - Average Atomic Mass

• The average atomic mass is the weighted average of all naturally occurring isotopes of the element.

Electron Shells

• Electrons exist in shells around the nucleus.
• The first shell has 2e, the second shell has 8e and the third shell has 18e
• Only electrons in the outer shell move to make bonds
• Each element has a different amount of electrons in the outer shell

Valence Shells

• Electrons are a type of glue that makes atoms bond together to make molecules
• Electrons in atoms can be classified as either outer or inner shell
• Bonding only involves the outer (valence) shell electrons

Nobel Gases

• Elements with fill outer or valence shells are in column 8 of the periodic table.
• Elements in column 8 are called inert gases because these elements do not combine or form bonds.
• All have 8 electrons in the valence shell except helium.

Non Metals - Gain Electrons

• For non metals electrons move to make the outer shell completely filled like the Nobel gases
• Elements can gain or lose electrons achieve the same number of 8 electrons in the valence shell
• Chlorine gains one electron to become like argon. Chlorine then becomes negatively charged

Metals - Lose Electrons

• Magnesium has two electrons in the valence shell.
• If it loses two electrons it has a full shell like argon.
• If magnesium loses two electrons it becomes doubly positively charged.

Ions

• An ion has a net positive or negative charge
• Charged particles that are charged as a result of a loss or gain of electrons
• Two types Lons, positive and negative lons
• The loss of electrons create Cations that are postive (+)
• The Gain of elctrons created Anions that are negative (-)

Chemical Bonds

• Two main types of chemical bonds including:
  • Ionic Compounds - Positive and negative charges join to form a bond.
  • Covalent bonds - Electrons are shared to achieve a complete valence shell.

Ionic bonds

• Formed between a metal and non-metal elements.
• Metals will transfer electrons to non-metals.

Net Charge

• Net charge must be neutral when two elements bind to each other

Anions and Cations

• Anions are particles the gain one or more electrions leading to an overall negative charge
• Cations are particles that lose one or more electrons, leaving them with a positive overall charge

Covalent and Lewis Dot Structures

• Helpful to understand covalent bonds.
• Nonmetals form mostly covalent bonds.
• Covalent bonds are when electrons are shared between atoms to achieve a total of 8 bonding electrons.
• Each covalent bond has two sharing electrons.
• Each atom shares the same number of electrons.

Lewis Dot Structures

• The electron dot structure consists of the element's symbol representing the nucleus.
• The element is surrounded by dots representing the element's valence electrons.
• Each dot represents an electron, and the number of outer shell electrons (valence electrons) can be located on the periodic table.
• The group number is the number of outer shell electrons.

Covalent Bonds of Carbon

• Carbon shares electrons with hydrogen.
• One electron shared by each set of atoms is a covalent bond and is represented by a solid line.
• Four covalent bonds means 8 valence electrons for carbon.

Types of Carbon Covalent Bonds

• Carbon can also make covalent bonds with other carbons
• Very unique property
• Classified as single, double and triple bonds

Electronegativity

• A covalent bond is a sharing of electrons between two atoms.
• When the atoms are shared equally, the covalent bond is even.
• Issues occur when atoms are defferent because it means not all electrons are shared equally

Electronegativity

• Unequal sharing means that the electrons will spend more time with a single atom.
• This makes the atom slightly negative in charge, where the atom on the bond will become partially positive.
• Bonds have partial negative and positive charges at each end, and covalent bonds are now classified as polar.

Electronegativity Tables

• Higher the value the stronger the pull of the atom type for the electron.
• Creates a polar covalent bond.
• Fluorine has the strongest pull for electrons
• Francium has the weakest attraction

Hydrogen Bonds

• The polar covalent bond has created one of the most important bonds in biology.
• It is also called the hydrogen bond.
• Oxygen in water has a slightly negative charge.
• Two hydrogens have a slightly positive charge.
• In effect the water molecule is attracted to other water molecules by the partial charges on the atoms.
• It is classified as intramolecular bonding.
• This attraction makes water a liquid with a boiling point of 100°C.

Learning Outcomes

• Identify the main groups of elements in the Periodic Table based on the atomic number.
• Calculate the atomic number based on the % of different isotopes.
• Calculate the number of electrons in the valence shell.
• Predict the type of bonding using the Lewis dot structures.
• Draw the molecular formula for ionic and covalent molecules.
• Describe how electronegativity creates hydrogen bonds.