Chemistry-Atoms

Atoms are the basic units of matter and consist of three primary subatomic particles: protons, neutrons, and electrons.

matter and atoms

matter

• a substance that has volume (occupies space) and has mass

• All matter is made up of atoms, which bond together to produce different substances

• Matter and energy are not the same thing

• Everything is either matter or energy

• Energy does not have any atoms

• Energy and matter can change into each other and vice versa

• The first element to be found in the world

atoms

• atom: building block of matter

• 1802: The first atomic theory of matter was presented by John Dalton. Dalton proposed that all matter is made up of tiny spherical particles, which are indivisible and indestructible

• We now know it is INCORRECT and atoms are made of smaller subatomic particles (protons, neutrons and electrons)

Pure substances VS Mixtures

• Matter can be split into two groups: pure substances and mixtures

Mixtures

• Mixtures: consist of 2 or more types of particles that are not chemically combined. Can be made of elements, compounds or both

  • air: oxygen gas (O2), Carbon dioxide gas (CO2), Nitrogen gas (N2)

  • Oil-water mixture

• Mixtures can be separated by physical means (e.g. filtration)

Pure Substances

• Pure substances: A pure substance is a substance made up of only one type of particle throughout. This means it has fixed composition and consistent properties. This can either be one single element or one single compound, but every sample of this substance that you examine must contain the same thing with a fixed, definite set of properties

  • pure element: copper metal (cu)

  • Pure compound: carbon dioxide gas (CO2)

• It cannot be separated by physical means (e.g. filtration)

Elements

• Elements: Made of just one type of atom

• It could be monatomic (i.e exist as individual atoms) or could also form molecules

• molecules: 2 or more atoms that are held together by chemical bonds

• elements cannot be separated into simpler substances by physical or chemical means

Compound

• Compound: Different types of atoms could combine to form new substances

• To determine whether a pure substance is an element or a compound, you must decide if the substance can be broken down into simpler substances

• Molecules: 2 or more atoms that are covalently bonded

• MOLECULES MUST BE COVALENTLY BONDED

• All molecules are compounds, but not all compounds are molecules

• Compounds are ionically bonded

Atomic structure

• An atom is made up of three types of subatomic particles:

  • protons (positively charged)

  • Neutrons (neutral)

  • Electrons (negatively charged)

• In chemistry, the word particle is a general term that refers to a small unit of matter

• Depending on the context, “particle“ could mean an atom, a molecule, an ion, or something else.

Atomic structures

• Protons and neutrons have relatively equal mass and are huge in comparison to electrons. This is why the nucleus is so dense

• Proton and neutron contribute nearly all the mass of the atom

• Protons and neutrons are relatively equally massive and are huge in comparison to electrons. This is why the nucleus is dense

• Proton and neutron contribute nearly all the mass of the atom

The Rutherford Experiment

• Rutherford’s model proposed the following

  • Most of the mass of an atom and all of the positive charge must be located in a tiny central region called the nucleus

  • Most of the volume of an atom is space, occupied only by electrons

  • The electrons move in a circular orbit around the nucleus

  • The force of the attraction between the positive nucleus and the negative electrons is electrostatic.

• Before Rutherford’s experiment, the atom was thought to be a spherical cloud filled with protons and electrons all over (think raisins caked, or plum pudding)

• In 1911, Rutherford’s experiment proved that the atom has a tiny but heavy nucleus and that most of the volume of an atom is space occupied by electrons

Bohr model

• In 1913, Niels Bohr developed a new model of the hydrogen atom that explained emission spectra. The Bohr model proposed the following

  • electrons revolve around the nucleus in fixed, circular orbits

  • The electron’s orbits correspond to specific energy levels in the atom

  • Electrons can only occupy fixed energy levels and

• Scientists quickly extended Bohr’s model of the hydrogen atom to other atoms

• They proposed that electrons were grouped in different energy levels, called electron shells

• These electron shells are labelled with the number n = 1, 2, 4

  The electron shells

• The highest energy level found is n7

• The highest energy level is found farther away from the nucleus because of the electrostatic attraction force

• The higher up you go, the higher the electrostatic attraction; therefore lower the energy, the closer you get to the nucleus

• If 2 negatively charged electrons share an electron shell have anti-clockwise spins so they never meet or come in contact with one another

• think 2 siblings circling each other

Different types of atoms

• The type of atom that makes up each element is determined by the number of protons (atomic number) in the nucleus

  • atomic number: the number of protons in the nucleus of the atom

  • mass number: the total number of protons + neutrons in the nucleus

  • Atoms are electrically neutral; therefore number of electrons = the number of protons

• The number of neutrons is calculated by: atomic mass - atomic number (number of electrons/protons)

• The atomic number is usually the smaller number, and the atomic mass is usually larger and has a lot of decimals

Isotopes

• all atoms that belong to the same element have the same number of protons in the nucleus and therefore the same atomic number,

• Atoms that have the same number of protons (atomic number) but different numbers of neutrons (and therefore different mass numbers) are known as isotopes

• Isotopes have identical chemical properties but different physical properties, such as mass and density. In particular, some isotopes are radioactive

• In nature, different elements have different numbers of isotopes. Gold only has one isotope, whereas lead has four isotopes, and mercury has seven isotopes

• Protons: determine elemental identity

• Electrons: determine chemical reactivity(how easily (or not easily) it is for the atom to undergo a chemical reaction )

• Neutrons: determine physical properties the physical properties of an element

Practice with Isotopes

• Identify the most abundant form of carbon out of the 3

1. ¹²₆ C - Carbon 12

2. ¹³₆ C - Carbon 13

3. ¹⁴₆ C - Carbon 14

• Carbon 12 - it appears in the periodic table, meaning it is the most commonly found Isotope

• Note that 12, 13, and 14 are the atomic masses of the atom

Electron configuration

• Using the Bohr model, it is possible to determine the basic electronic configuration of atoms by applying the following rules

1. Each shell can only contain a maximum number of electrons

2. Lower energy shells are full before higher energy shells

Steps to annotating an atom - Electron configuration

1. Determine the number of electrons in your element

2. Recall the maximum number of electrons each shell can hold

3. Place the number of electrons in the shells from the lowest energy to the highest energy. Do not exceed the maximum number of electrons allowed

4. Write the electronic configuration by listing the number of electrons in each shell separated by commas. (2, 8, 1) (for sodium)

• THIS ONLY APPLIES TO THE FIRST 18 ELEMENTS (TILL ARGON)

Atoms

• Atoms: Building block of matter

Ions

• A positively or negatively charged atom or group of atoms

• Remember that an atom has the same number of electrons and protons, which means that it is neutral

• When atoms pick up an additional electron (s) or lose electrons (s), there is no longer a balance between the positive and negative charges

• They lose electrons/protons to achieve a full valence shell (octet rule)

  • An exception to the octet rule would be the duet rule (applies to hydrogen & helium)

  • The duet rule states that there would be 2 electrons in the valence shell

The 2 main types of ions

1. monoatomic, e.g.

   a. Na⁺, Cl^-, Cu⁺²

2. Polyatomic

   a. e.g. pos₄^3-, OH^-

Cations

• a positively charged ion (an atom loses electrons)

  • Meg2+, Al3+

• Magnesium, for example, has 2,8,2

• It loses 2 electrons and becomes

  • Mg ^+2/2+

  • You would also put square brackets around the visual representation and put a 2+/+2 in the upper right corner

Anions

• Negatively charged ion (atom gains electrons)

  • e.g. Cl-, O2-

• Chlorine, for example, has 2, 8, 7

• It gains electrons and becomes

  • Cl^-

  • There is an implied one

Metals

• Metals have loosely held outer shells (valence electrons)

• They can lose these electrons to become cations, this is because the resulting ion has fewer electrons than protons

Nonmetals

• Non-metals attract electrons

Transition metals

• They can take on more than 1 charge

• They have variable charges

• Really big valence shells so they can lose or gain electrons easily

positive control: variable with known result (e.g. plant is exposed to sunlight)

negative control: absence of IV (e.g. plant exposed to no light)

• No result (e.g. no change in plant height/mass)

Electron Transfer and Ionic Bonding

Recall classifying matter

atom: the smallest particle of matter

Element: a substance made from only one type of atom

Compound: A substance made from 2 or more different types of atoms

Molecules: A substance made from two or more atoms that have chemically combined

Ionic compound

• An ionic compound typically forms when a metal reacts with a nonmetal

• Ionic compound = metallic cation + non metalli anion

• During a reaction, there is a transfer of electrons. Metals → Non-metal

• Once this occurs, the oppositely charged ions join together in a lattice

• Oppositely charged ions are attracted to one another, and this attraction would be called an “electrostatic attraction”

• electrostatic attraction, AKA ionic bonds, AKA the attraction between oppositely charged ions

• They are brittle, as when a force is applied to the lattice, the ions with like charges align and actively repel each other, shattering the lattice

• The ionic bond is very strong and brittle because rearranging ions causes repulsion between like charges

• You are sliding the layers so the bond is still intact, but the layers slide and the negatives end up next to negatives, so they repel each other, and then the lattice breaks apart

• They normally have very high mpts

• Energy must be transferred to a substance to make it melt or boil

• This energy overcomes the strong electrostatic forces of attraction, which act in all directions between the oppositely charged ions

  •  Some forces are overcome during melting

  • All remaining forces are overcome during boiling

• The more energy needed, the higher the melting point or boiling point

• Since the electrostatic forces of attraction between oppositely charged ions are strong, their melting and boiling points are high

• Solid sodium chloride needs a temperature of 801 degrees to melt it

Electron transfer diagrams

• Electron transfer diagrams are used to show the path that electrons take when they are removed from a metal and added to a non-metal during ionic bonding

• show the transfer of electrons from metals to non-metals

Steps to draw an electron transfer diagram

1. Draw an electron shell diagram of the neutral metal and non-metal

2. Add a ‘+‘ between them

3. Draw an arrow leading from each valence in the metal to the valence shell of the non-metal

4. Add an arrow towards the resulting ions

5. Draw an electron shell diagram of the resulting cation/s and anion/s

6. Write the chemical symbol of the metal and the non-metal in the centre of the electron shell diagram, taking care of the electron shell diagram, taking care to add charges to your ions.

Naming Simple Ionic Compounds

RULES:

1. Name the cation first (metal) before the anions (non-metals)

   *If metal is a f metal is a transition metal, indicate the valency in numerals after the name
   E.g. Fe(III), Ag(I), Gold (I)

2. The name of the cation remains as is

   E.g., Sodium ion, Na⁺ ion

3. The name of the non-metal anion is changed. Its suffix becomes ‘-ide’
   

   

Simple Ionic Compound Formulas

Metallic bonds

• consists of the attraction between positively charged metal ions and delocalised electrons, allowing for malleability and conductivity. (can only be cations)

• They tend to lose their outer shell electrons easily and turn them into positively charged cations.

• 

• the electrostatic force of attraction between the positively charged metal cations and negatively charged valence electrons occur in all

Alloys

• a mixture of two or more metals, combined by metallic bonding

• they tend to harder than pure metals because pure metal atoms being the same and arranged in layers , as opposed to alloys that contain element with atoms of different sizes

common alloys:

• Steel: iron(metal)+ carbon(non-metal)

• Bronze: copper(metal)+ tin(metal)

• Brass: copper(metal)+ zink(metal)

Types of bonds