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Exclusively Yours,

Preliminary HSC Chemistry Syllabus Notes

Module 1

Properties and Structures of Matter

Week 1 Material – Properties of Matter

(Version 1)*

*Notes will consistently be revised through the year to ensure quality. 

Week 1 Content

Overview of Week 1’s Inquiry Question

Learning Objective #1 – Homogenous vs Heterogenous Mixtures 

Learning Objective #2 – Separation techniques based on physical properties

Learning Objective #3 – Applying IUPAC method to name inorganic substances 

Learning Objective #4 – Chemical, physical property and position of elements’ trends on the periodic table 

HSC Biology Syllabus Lecture Video – Properties of Matter

Week 1 Homework Set (Essential for Band 5)

Curveball Questions (Moving from Band 5 to Band 6!)

Solutions to Week 1 Homework Set

Course Theory

Before we hop on the materialistic train and start digging into the content, please give me a minute to walk you through what you should keep in mind as the ‘major highlights’ for this week’s material. 

The inquiry (overarching) question for this week deals with classifying and separating substances based on their properties (chemical and physical). 

Under the concept of classifying substances, we explore the concepts of homogenous and heterogenous mixtures, metal, semi-metals and non-metals. 

Following this, we will learn how to name inorganic substance according to their systematic names.

To finish off this week’s notes, we will have a look at how elements’ position on the periodic table is based on their physical and chemical properties

What is an element? 

Elements are single atom species (‘particles’). 

They are just one atom, not combined or bonded to another atom. Each box on the periodic table represents an unique element.

These elements (single atom) can bind with other or same the elements (single atom) to make up all things in the universe! This includes your cells. 

If you do not know what cells are, your body organs, your skin are all made up of millions of cells. There is about 100 trillion atoms in a cell. 

Everything from your schoolbag to trees are made up of atoms. Literally everything. Why do things look different then? Why does a tree doesn’t look like a human? 

Different chemical compositions (atoms), arrangement and the various chemical bonds between atoms (elements).

Every neutral atom for any given element has at least one proton, neutron and electron. Depending on the element, they will have a different amount. 

You can find this amount by using the periodic table. I will show you how now.

          I      II                                                                                   III   IV   V    VI   VII  VIII

Take the element: C (top right). C represents the element Carbon. 

Take another element: Cl (top right). Cl represents the element chlorine.

There is a number ‘6’ above C. This is called the proton number. It represents the amount of proton a neutral carbon atom would have. The number below ‘C’ is the atomic mass number for carbon. Here, the number is 12.011. This number represents the total amount of proton and neutrons combined. For a neutral element’s atom (any box on the periodic table), the proton number also represents the amount of electrons! Hence, carbon would have 6 protons, 6 neutrons (~ 12.011 – 6) and 6 electrons!

 

Molecules are species that are made up of two or more atoms bonded together!

Compounds are species that are made up of two or more elements’ atoms’ bonded together!

All compounds are molecules but not all molecules are compounds (e.g. oxygen gas, O2, is not a compound but is a molecule). 

Ions are charged atoms! An ion can either be positively or negatively charged. Positively charged ions are called cations. Negatively charged ions are called anions. 

If a neutral atom loses an electron (negatively charged), they will become a cation. If a neutron atom gains an electron, it will become an anion! 

Now, we have some brief understanding of elements and atoms and the sort of substances that they can form, let’s move to explore some of this Week’s learning objectives!

Learning Objective: Explore homogenous and heterogenous mixtures through practical investigations

 Using separation techniques based on physical properties!

This learning objective requires us to separate homogenous and heterogenous mixtures based on their physical properties

What are homogenous and heterogenous mixtures?

Mixtures is the combination of two or more substances. Because of this, mixtures may have a variety of physical and chemical properties, depending on the substances that make up the mixture. Some common categories of mixtures include: Solutions, Suspensions, Colloids and Emulsions.

Let’s explore each at a time!

Solutions

A solution has two components. These are solute and solvent. 

Solutions are homogenous mixtures. Homogenous means that the solution has a CONSISTENT appearance and chemical composition throughout the mixture. 

‘Homo’ means same or uniform.

‘Hetero’ means different or not uniform.

An example of a HOMOGENOUS mixture would be salt water. Inside salt water, there are sodium and chloride ions (charged atoms) that are UNIFORMLY (evenly) dissolved in the water. Because of this, when you look at a glass of salt water, you will not see any solids or ‘unevenness’ 

Going back to the topic of solutions, we said that a solution is made up of a solute and a solvent. Solute(s) is the substance that is dissolved in the solvent. The solvent is the medium which the solute is dissolved in. 

For example, if you throw a teabag of sugar into a cup of water, the sugar will be dissolved in the water (provided that there is enough water for complete dissolution). The sugar will then be the solute and the water would be the solvent. Collectively, the solute and solvent makes up the sugar water SOLUTION (homogenous mixture). 

Colloids

Colloids are another example of a homogenous mixture. Night cream is a colloid. This is because inside the cream, there are large chemical molecules throughout the mixture as opposed to small molecules in a homogenous mixture such as a solution. Regardless of the size, since the chemical composition of colloids are uniformly throughout the mixture, you will NOT see any ‘unevenness’ in the appearance.Therefore, colloids are homogenous mixtures.

The difference between colloids and solutions:

Due to the large molecules in colloids, most colloids are much less transparent in appearance than solutions due to the effect of light scattering when light hits the large colloids particles in the mixture. Light scattering is significantly less in solutions (smaller particles) compared to in colloids. This is why you may not be able to see through your night cream.

Now that we have discussed briefly about homogenous mixtures, it is time to venture out to explore the world of heterogenous mixtures!

Heterogenous Mixtures

Definition wise, heterogenous mixtures are mixtures that are not uniform in their chemical composition and physical appearance.

Usually, there are more than one observable state of matter (gas, liquid, solid) in a heterogenous mixture.

An example of a heterogenous mixture is dirt at the bottom of a bucket of water. The dirt is not uniformly mixed in the mixture if you gives a little bit of time to allow the dirt and water to settle.

Suspensions

Suspensions are heterogenous mixtures. The dirt particles at the bottom of the bucket of water is an example of a suspension. Suspensions have undissolved solids in a liquid medium. 

Emulsions

Emulsions is a heterogenous mixture made up of two liquids. A typical example of an emulsion is milk. Water is mixed throughout the milk BUT not dissolved. This is why you sometimes see water droplets near the top of the container’s inner walls. Those water droplets rose above the milk. This observable inconsistency in appearance is why emulsions are considered heterogenous mixtures. Another example is a mixture of oil and water.

NOTE: There is no such thing as a heterogenous solution ONLY a heterogenous mixture. A solution is homogenous if it’s physical appearance and chemical composition are uniform. If you write ‘heterogenous solution’, you may risk not getting a mark in the exam.

We have now touched on the differences between homogenous and heterogenous mixtures as well as an example of each! 

It is time to figure out how to separate these homogenous and heterogenous mixtures using their physical properties!

But what are physical properties?

Physical properties are any features that is observable and can be tested without altering the substances’s chemical composition! It may or may not involve a physical change.

Testing boiling and melting points for instance involves a physical change (altering the states of matter such as from solid to gas). However, the species is not being converted into another identity is not altered. For example, if you boil liquid water to produce water vapour, in both cases, the species is still water. Only it is a physical change. Not a chemical change. 

Chemical properties are any features that can be observed and tested that involve altering the substances’s chemical composition. That is, it involves a chemical change. Some examples of tests for chemical properties include:

Chemical bond reactivity, heat of combustion, toxicity, pH (pH tests by the substances reacting with ions to change colour)

Separation techniques used on mixtures based on their physical properties!

Simple Separation – Sorting heterogenous solid mixtures by colour. For example, sorting coloured marbles (solids). 

Magnetic Separation – Sorting solids because on their magnetic properties (magnetic vs not magnetic) in a mixture

Decanting – Separating a precipitate (solid) from liquid. This can be done by pouring the liquid from the container, leaving the solid behind. 

Sedimentation & Flotation – Separating mixtures by density. This is usually used for mixtures containing both liquid and (insoluble) solid.

Separating funnel – Separating immiscible (cannot be mixed) liquids by density (Higher density liquid will be at the bottom)

Centrifuging – Separating solids by density (Higher density particles will be at the bottom)

Evaporation & Crystallisation / Distillation & Condensation – These four separation techniques based are based on melting/boiling/freezing point of components in the mixture. Evaporation is based on differing boiling points. The substances in the mixture with the lowest boiling melt will turn into gas first. It is not uncommon for distillation to be used in conjunction with evaporation so that the gases are collected. Crystallisation is a measure of freezing point. The components of a mixture with the lowest freezing point will crystalise (turn into solid) first. Condensation is the process of a gas turning (condensing) into a liquid. Generally, substances with higher boiling point will have a higher freezing point.

Solubility – Dissolving a mixture in a solvent to separate it into its components. Different components in the mixture needs to have different solubilities.

Paper Chromatography – A type of solubility separation technique. For example, mixtures of different dyes are dissolved in water, such as a wet chromatographic paper that is hung upright with one end soaked in the dye mixture. The dye molecules in the mixture will move up the paper and be separated according to the solubility of different dyes. The most soluble dye will climb the highest (we see this because each type of dye will need to have a different colour). The result is a pattern of colours along the chromatographic paper, each colour representing a different dye in the mixture. 

Now, suppose that we have separated a mixture into its individual components!

What should we do if we want to know how much of each component contribute to the overall mixture?

One method is to calculate the percentage composition by weight of each component element and/or compound in the mixture which we have separated!

Learning Objective: Calculating percentage composition by weight of component elements and/or compounds

Below is a general rule for determining the % mass for each element in a compound

The formula for composition by mass (% mass) of a substance = element’s mass on periodic table / total molar mass of compound

Molar mass = mass per mole. You can find this on the periodic table 

Number of moles (n) = element’s mass (g)/molar mass (g/mol). This can also be expressed as n = m/MM

Sometimes, you see molar mass written as M instead of MM. However, the unit of concentration (molarity) is also M. So, i prefer using MM for molar mass.

Avogadro’s number = 6.022 x 10²³

Number of moles = number of particles / Avogadro’s number

Gravimetric Analysis

Gravimetric analysis is a technique used to find the percentage composition of an element or compound in a mixture:

1. Weigh a compound that you want to determine its percentage composition by mass of its different element components

2. Dissolve this compound completely in water, forming ions.

3. Select and add an appropriate solution to the compound so that the ions of the compound will form a precipitate (solid). It is important for you to add EXCESS of this appropriate solution so to ensure that all ions (of the compound) are precipitated out of solution. 

4. Gently heat the mixture so that only the precipitate is left behind (you can choose to filter the solution before heating too). Stop heating when the mass of precipitate is constant. 

5. Weigh the precipitate

6. Work out the moles of the precipitate (n=m/MM) and use the moles of the precipitate to determine the mole of each element in the compound via molar ratio

7. Multiply elements’ moles by their molar mass to obtain each element’s mass

8. Divide each element’s mass by the compound’s total molar mass then multiply by 100 to obtain the % composition by mass for each element in the compound.

Gravimetric analysis is use din many analytical work. For example, it can allow forensic chemists to work out the chemical composition of some  substance found at a crime scene. 

 

We have explored the definition of compounds. We will now move on to learn how to name inorganic compounds!

What is the difference between organic and inorganic compounds? Organic compounds ALL contain carbon atoms (often attached to hydrogen atoms). Inorganic compounds mostly deal with compounds without carbon atoms. However, in some cases, inorganic compounds may have carbon atoms. There is no 100% clear distinction.

Learning Objective: Investigate the nomenclature of inorganic substances using the IUPAC naming conventions

Nomenclature is the way of naming species in chemistry. This learning requires us to name inorganic substances. 

The types of inorganic substances that we will be exploring are ionic compounds, molecular compounds and acids.

Ionic compounds are substances made up of a metal and a non-metal. Using the IUPAC naming convention for ionic compounds, we first name the metal followed by the root name of the non-metal and then add ‘ide’ to the end of the non-metal’s root name.

 

Ionic Compounds

Ionic compounds are substances that are made up of a metal and a non-metal. Using the IUPAC naming convention for ionic compounds, we first name the metal followed by the root name of the non-metal then add ‘ide’ to the end of the non-metal’s root name. 

Example: How do you name the ionic compound, NaCl?

1. Identify the metal. In our example, this will be Na (Sodium)

2. Identify the non-metal. In our example, this will be Cl (Chlorine)

If you are not sure of how to identify whether something is a metal or non-metal, we will cover this in the next learning objective. So bare with me for now.

3. Identify the root name for the non-metal, in this case it will be ‘chlor’ for chlorine. You will need to get used to the root names of non-metals. The more practice questions you do, the more root names you will get exposed to. For instance, oxygen has the root name ‘ox’. Iodine has the root ‘iod’, etc. 

4. Add ‘ide’ to the end of the non-metal’s root name. So, root name for chlorine is ‘chlor’, after adding ‘ide’, this will become chloride

5. Write out the full name for the ionic compound. For our example, this will be, Sodium Chloride! 

 

Common root names for non-metals

By the end your HSC, your should know more root names than what is on this list!

Nitrogen (N) = Nitr

Phosphorus (P) = Phosph

Oxygen (O) = Ox

Iodine (I) = Iod

Fluorine (F) = Fluor

Chlorine (Cl) = Chlor

Bromine (Br) = Brom

Sulfur (S) = Sulf

Hydrogen (H) = hydr 

Carbon (C) = carb

Covalent Molecular Compounds (Non-metal with Non-metal)

How do you name these compounds with non-metal bonding with other non-metal?

How would you name N2O5? 

1. First, name the non-metal that is closest to the top left of the periodic table. Here, this will be the nitrogen element (not oxygen)

2. Add a prefix to the non-metal’s name (more on these prefixes later) to the non-metal if there is more than one atom in the compound. For example, in N2O5, there are two nitrogen atoms. 

3. Name the root name of the second non-metal (like we did with ionic compounds). Here, it is ox for oxygen.

4. Add ‘ide’ to the end of the second non-metal’s root name (like we did with ionic compounds). This will be oxide for oxygen.

5. Add a prefix to the root name of the second non-metal if it has more than one atom (e.g. in N2O5,  there are five oxygens).

6. Write out the name! Dinitrogen Pentoxide.

NOTE: The ‘Di’ and ‘Pent’ in front of nitrogen and oxide are prefixes! 

 

PREFIXES FOR NAMING COVALENT MOLECULAR COMPOUNDS

Exceptions to the above prefixes rule: If the non-metal starts with an “a” OR “o” (e.g. oxygen) and the prefix used also ends with either an “a” OR “o”, we remove the “a” or “o” at the end of the prefix. BOTH conditions (for the non-metal side) and (prefix side) must be satisfied in order to such exception (removal) to be applied.

Example of exception: N₂O₄. The name is NOT Dinitrogen TETRAoxide. It is Dinitrogen Tetroxide (prefix is Tetr instead of Tetra) In this case, the non-metal (oxygen) starts with a “o” and the prefix ended with an “a” (2 conditions satisfied). So, we remove the “a” from Tetra. 

Another example of exception: CO. The name is NOT carbon monooxide. It is carbon monoxide (prefix = mon instead of mono). 

Naming acids WITHOUT oxygen atoms

1. Write ‘Hydro’

2. Find the root name of the non-metal

3. Add ‘ic’ to the end of the non-metal’s (or metal) root name

4. Write ‘acid’

Example: Name the acid HCl. Notice that there is NO oxygen atoms in the acid, HCl. So we can name it as follows:

Hydro + Chlor + ic + acid     (all 4 steps in order)

Hence, full name for HCl is Hydrochloric acid. 

Naming acids WITH oxygen atoms

1. Ignore the hydrogen atoms. Do not write hydro.

2. Find the root name of the non-metal (or sometimes transitional meta).

3. Add ‘ic’ to the end of the non-metal’s root name.

4. Write ‘acid’

Example: Name H₂CrO₄

Ignoring the hydrogen atom, we find that the non-metal or transition metal. In this case it is the Cr (chromium) atom. The root name for the chromate is chrom. 

Since, we add ‘ric’ to the root name ‘Chrom’, making it chromic. 

Lastly, we add ‘acid’ to the end. 

Therefore, HCrO is called Chromic acid.

 

Exceptions to acids containing oxygen are those acid with sulfate and phosphate ions!

Example: H₂SO = Sulfuric acid (NOT sulfic acid). Need to add ‘ur’ to the root name ‘sulf’

Example: H₃PO = Phosphoric acid (NOT phosphic acid). Need to add ‘or’ to the root name ‘Phosph’

 

Common polyatomic ions (You need to know more than what is on this list. As you do questions, you will come into contact of more of these ions!)

Carbonate – CO

Nitrate – NO

Permanganate – MnO

Phosphate – PO –

Sulfate – SO³⁻

Chromate – CrO²⁻

Cyanide – CN

If the above common polyatomic ion version has one less oxygen, the ‘ate’ is reduced to ‘ite’. See below.

Sulfite – SO² (compared to Sulfate – SO²⁻)

Phosphite – PO³⁻ (compared to Phosphate – PO³⁻)

Chromite – CrO²⁻

NOTE: These common polyatomic ions with one less oxygen has effect on the naming of oxygen-containing acids!

Recall that the name HPO (with phosphate ions) is called Phosphoric acid

Well, if the phosphate ions has one less oxygen atom, HPO (with phosphite ions instead of phosphate), it is now called Phosphorous acid. 

The difference is that instead of adding ‘ic’ to the end of the ion’s root name, we add ‘ous’ instead.

Another example: HNO is called Nitric acid. However, if the nitrate ions has one less oxygen atom (nitrite ion) and it is present in the form, HNO, it is now called nitrous acid (rather than nitric acid)! 

 

So, here is a general rule of thumb:

Remember how common polyatomic ions (nitrate, phosphate, sulfate, etc) in terms of how many oxygen atoms they have.

This way, if you see a polyatomic ion that has one less oxygen atom than its common version, you will know.

Having one less oxygen atom will affect the name of the acid. Acids containing polyatomic ions with one less oxygen than the common version of the ion will end with ‘ous’ acid. (e.g. nitrous acid) instead of the ‘ic’ acid (e.g. nitric acid). You can imagine that the ‘ous’ acids are the ‘weirder’ versions of the ‘ic’ acids. 

Okay! One last learning objective before we wrap up this week’s notes. Well done for making it this far!

Learning Objective: Classify the elements based on their properties and position on the periodic table through their:

Physical properties

Chemical properties

Please note that these trends that we will be exploring are general trends. They do not apply successfully for every element on the periodic table. Real life is never so simple it is. Haha! This is why life is interesting. Anyways, let’s get in to have a look at the GENERAL TRENDS of the periodic table!

Trends for elements’ position on periodic table based on physical properties

Atomic Radius

The size of an element’s atom’s radius is determined by how far its electrons are from the centre of the atom (nucleus). You may be aware already that electrons exist in energy shells.

The more energy shells an element’s atom has, the larger its radius will be. You also may know that energy energy shell can only hold a maximum amount of electrons. The energy shells that are further away from the atom’s nucleus can hold more electrons than the energy shells that are closer to the nucleus.

The 2(n)² rule can be used to give you an idea of how many electrons an energy shell can hold, where n = energy shell number.

Energy Shell 1 (closest to nucleus): 2(1)² = 2

Energy Shell 2 (2nd closest to nucleus): 2(2)² = 8

Energy Shell 3: 2(3)² = 18

Energy Shell 4: 2(4)² = 32

so on, so forth.

You may be aware that protons and electrons attract (opposite charges attract). Protons are positively charged and electrons are negatively charged. An atom’s nucleus consists of protons and neutrons (uncharged/neutral) with electrons orbiting around the nucleus. The stronger the effective nuclear charge (or the attractive force pulling electrons towards the nucleus), the smaller the atomic radius will be. This is because electrons will be closer to the nucleus. Recall that atomic radius is determined by the distance between the outermost electron (valence electrons) and the centre of an atom’s nucleus.

NOW with these fundamentals revised, we will proceed back to the trends on the periodic table!

As you move from left to right across any given row (period) on the periodic table, the size of the elements decrease. By size i mean the atom’s radius for an element. 

Reason: Atomic radius decreases as we move from left to right on any given period because proton number increases as we move from left to right! It increases by one for every box you move towards the right. In case you are wondering, proton number represents the number of protons a certain element’s atom has. Yes, indeed, as we move from left to right, the number of electrons also increase proportionally (1:1) with the number of protons. 

However, the extra electron is put in the same energy shell (remember energy shells can hold more than one electron!) as the other elements on the same row on the periodic table. This allows the effective nuclear charge to have a much greater effect than if the extra electron were put in a new energy shell that is farther away from the nucleus (as we move from left to right on a given period)

As we move down any column (group) on the periodic table, the size of atomic radius increases. This is because for every box you move down the periodic table, a new energy shell is created. 

Simplifying it to an one proton to one electron basis: Suppose there are two atoms. Atom A and Atom B.

Suppose you add one proton and one electron to Atom A. Similarly, you add one proton and one electron to Atom B. Both Atom A and Atom B has the same number of protons, neutrons and electrons. However, the new electron that you added to Atom A is put to a new energy shell compared to the new electron on Atom B which is on the same energy shell. What is the atomic radius for the Atom A and Atom B?

Well, Atom A is going to be bigger than Atom B (in terms of atomic radius). This is because a NEW energy shell (that is the furtherest away from the atom nucleus) is added to Atom A. Remember, both Atom A and Atom B has the same number of protons, neutrons and electrons. The effective nucleus charge experienced by the electrons in Atom A will be weaker than in Atom B. This is because of the concept of ‘electron shielding’. This will be covered in later week’s notes with greater detail. But, for now, electron shielding means that energy shells closer to the nucleus will act as a shield (against the effective/positive nucleus charge exerted by the protons) for the electrons in energy shells that are further away from the nucleus. So, energy shell 1 (closest to nucleus) will act as a shield for energy shell 2. Energy shell 2 will act as a shield for energy shield 3, etc. This means the more energy shields an atom has, the less effective nucleus charge is felt. Hence, the smaller the atomic radius is going to be. 

Atomic Mass

Atomic mass increases as you move towards the bottom right of the periodic table. This is because as you move from left to right, the number of proton, neutron and electrons increase. 

Metallic characteristics

As you move from left to right, the metallic characteristics of the elements decreases. Metallic characteristic can be interpreted as how readily an element is able to donate its electron. So, you may now see that as metals are on the left side of the periodic table and non-metals are on the right side of the periodic table. As effective nuclear charge decreases as you move down a given group on the periodic table. Metallic characteristics decrease as you move from left to right because elements are more readily accepting electrons (electronegativity increases) rather than giving up or donating electrons.

In between metals and non-metals (on the periodic table) are semi-metals also known as metalloids. These are Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Tellurium (Te) and Polonium (Po). These typically have properties that are a mix of both metal and non-metals. 

GENERALLY SPEAKING, Metalloids’ physical characteristics typically resembles of a metal, however, its chemical properties typically represent that of a non-metal. 

 

Some physical characteristics that differentiates metals and non-metals

Melting and Boiling Point

Melting and boiling point increases then decreases as you move from left to right on the periodic table.

As you move from metal towards (but not reaching) semi-metals , melting & boiling point increases. This is because metallic bonding increases in strength. More on metallic bonding in later weeks. Basically, since the number of electrons increases as you move across a period, metallic bonding increases in strength.

However, as you move from semi-metal to non-metals, melting/boiling point decrease because no metallic bonding exist in non-metals. This is because most non-metals exist as gases such as oxygen and fluorine at room temperature. They are held together by intermolecular forces (mainly dispersion forces) that is due to the interaction of electrons. 

Another factor that contributes to melting and boiling point are intermolecular forces. As the number of electrons increase, the stronger the dispersion will be. As you move down a given group on the periodic table, the number of electrons increases. Thus, melting and boiling point increases.

Furthermore, the more organised the atoms are, the stronger the dispersion forces will be. The arrangement of atoms decreases (becomes more random and scattered) as you move from metals to non-metals. Non-metals are commonly found as gases so they tend to ‘float’ around in space, held weakly together by intermolecular forces.

On the other hand, metals have a more organised structure, held together by metallic bonding.

Trends for elements’ position on periodic table based on chemical properties

Ionisation Energy

1st Ionisation Energy is the energy required to strip off a valence (outermost) electron from an atom, converting a neutral atom into a cation as a result.

An ion is a charged species. A cation is a positively charged ion. Since the atom lost an electron (a negative charge of -1), the overall charge of the ion will be +1

Ionisation energy increases as you move towards the top right of the periodic table. The reason why 1st ionisation energy increases from bottom to top is because the effective nucleus charge increases as you move up groups and across periods on the periodic table. The more loosely bound a valence electron is, the less energy is required to remove an electron from the atom. 

NOTE: 2nd Ionisation energy is greater than 1st ionisation energy as more energy is required to strip an electron away from a cation. Cations will be more strongly attracted to its electrons, hence, stripping an electron away from it will be require more energy. 

Furthermore, in some exam questions, you may come across that in order to remove an electron from an atom, it will cost SUBSTANTIAL ionisation energy compared to other elements. This is likely due to you are trying to remove an electron from an element with its valence (outermost) energy shell already COMPLETELY filled. Removing an electron from an element with an already-filled-valence shell will cost MUCH MORE energy than removing an electron situated in a valence shell that is not 100% filled. 100% filled valence shells are the most stable form of an atom! DONT TOUCH EM!

Electron Affinity & Electronegativity

Electron affinity is an atom’s change in energy when an electron is ADDED to the atom to become an anion (negatively charged ion) The units for electron affinity is typically in kilojoules per mole (kJ/mol). Electron affinity is determined by ionising an atom into its gaseous state. You do not need to worry about ionising atoms to test for electronegativity.

Electronegativity is the ability for n atom to attract a pair of electron. It is a unit-less measure. It is just a number assigned to an element relative to other elements.

Because nuclear effective charge is low for elements situated at the bottom of the periodic table, it is difficult for the nucleus to attract that an electron. Hence electronegativity decreases as you move towards the bottom of the periodic table. Furthermore, electronegativity as you move from left to right because the number of protons increase (and electron shielding effect is negligible when moving across a given period than going down a group), increasing electronegativity.

Furthermore, since the number of electron increases (by more than 1) as you move down a box for any given group, the extra number of electrons will create greater repulsion forces that lowers electronegativity.

The trend for electron affinity is the same for electronegativity (increases as you move towards the top right of the periodic table)

 

Chemical reactivity

An atom’s chemical reactivity is determined by its electrons. The more easily an atom is able to freely interact with other atoms (donate or share its electrons), the more reaction the atom will be. 

As electrons are loosely bounded to the nucleus, chemical reactivity increases as you move towards the bottom left of the periodic table. This is effective nuclear charge is the weakest at the bottom left of the periodic table. 

More ‘Advanced’ Nomenclature 

Recall that at the beginning of this week’s notes, there was a diagram of a periodic table with group numbers above it – Group I, II, III, IV, V, VI, VII and VIII.

Valency of atoms is important in assisting you to write chemical formulas for compounds!

Valency of atoms is essentially the combining power of the atom with a hydrogen atom (one electron) to form a compound. 

An atom with a valency of +1 means that it wants to donate one electron. An atom with a valency of -1 means that it wants to accept one electron. 

Metals wants to donate electrons and non-metals tends to accept electrons. Thus, metals have positive valencies and non-metals have negative valencies.

  • All elements under Group I have a valency of +1
  • All elements under Group II have a valency of +2 
  • All elements under Group III have a valency of +3.
  • All elements under Group IV has a valency of +4.
  • All elements under Group V has a valency of -3.
  • All elements under Group VI has a valency of -2
  • All elements under Group VII has a valency of -1.
  • All elements under Group VIII has a valency of 0 (zero combining power with hydrogen to form compounds. Group VIII elements are chemically inert (not chemically reactive by nature).

Most of the compounds that you will be dealing with throughout HSC Chemistry are going to be neutral. This means overall charge (‘valency’) of the compound must be zero. 

However, there can be charged compounds such as polyatomic ions and complex ions. We will touch on these in a later topic (oxidation states)

Example of neutral compound: NaCl (net charge = 0)

Example of charged compound: CrO₃²⁻ (net charge = negative 2)

So, let’s explore how you can write the chemical formula for some neutral compounds!

Example: Why is sodium chloride written as NaCl and not NaCl2, Na2Cl, Na4Cl, etc. This is because sodium is in Group I. It has a valency of +1. Chlorine is in Group VII. It has a valency of -1. Overall, the +1 and -1 ‘cancels out’ each other, making a net charge of zero. The NaCl compounds is neutral

You could technically write NaCl as Na2Cl2 (+2 and -2 cancels out to make zero net charge). However, according to IUPAC nomenclature, you will be wrong. HSC Chemistry follows IUPAC like we did earlier!

Another example: How to write Iron(III) chloride? Well, the (III) means that iron has a valency of 3+. Chlorine is in group VII. This means that chlorine has a valency of -1. We need three chlorine atoms to ‘balance out’ that +3 charge of that one Fe3+ ion. So, we write it as Iron (III) chloride as Fe(Cl)₃

We put brackets around chloride with a subscript ‘3’ to indicate that there are three chlorine atoms each individually bonded to the Fe (iron) atom.

So what does 2Fe(Cl)₃ mean? Well, the coefficient ‘2’ means that there are two Fe(Cl)₃ molecules! 

HSC Biology Syllabus Lecture Video – Properties of Matter

[Video will be uploaded HERE SOON!]

Week 1 Homework Set (Essential for Band 5)

Question 1: Distinguish the terms physical property and chemical property of an element [4 marks]

Question 2: Distinguish between ionisation energy and electronegativity [4 marks]

Question 3: Explain in detail four trends on the periodic table [8 marks]

Question 4: Describe the difference between a homogenous and heterogenous mixture [4 marks]

Question 5: Explain the purpose of gravimetric analysis [4 marks]

Question 6: Describe how would you perform gravimetric analysis for the NaCl compound [6 marks]

Question 7: Research the methodology behind all the separating techniques that was mentioned in this week’s notes under the heading “Separation techniques used on mixtures based on their physical properties!”

Question 8: A friend of yours said that “your rich-looking pizza is a homogenous mixture.” Do you agree or disagree? Explain.

Question 9: Name the following compounds, feel free to use the periodic table to look up the name of the elements according to their symbols

  • CaCl₂
  • CaF₂
  • MoS₂
  • FeCl₃
  • CuCO₃
  • CoCO₃
  • PbO
  • PbO₂

Question 10: Write out the chemical formula for the following compounds, feel free to use the periodic table to look up the symbols for the elements

  • Manganese (IV) oxide
  • Magnesium phosphate
  • Iron (III) oxide
  • Gold (III) chloride
  • Sulfuric Acid
  • Hydrogen peroxide
  • Ozone (feel free to google this one, haha)
  • Boric acid
  • Beryllium oxide
  • Antimony hydrate
  • Sodium hydrate
  • Copper (II) sulfide

Curveball Questions (Moving from Band 5 to Band 6!)

Curveball Question 1: A forensic scientist needs your help! He cannot figure out exactly how an unknown compounds look like (chemical composition wise). All he knows is that, in the compound, there is 14 grams of calcium, 11 grams of oxygen and 0.7 grams of hydrogen. He wants you find the simplest (empiricial) formula for the unknown compound that he is working worth. Can you help him?

Curveball Question 2: Another forensic scientist needs your help! He analysed a sample of cereal. He found that it contains 57% carbon, 27% oxygen, 9% nitrogen and 6% hydrogen. He wants you to find out what is the simplest formula for the compound that makes up the cereal. He also wants you to find out the molecular formula, all he knows is that the cereal has a molar mass of 294 grams per mole.

Solutions to Week 1 Homework Set

Solutions to each week’s homework set will be uploaded one week subsequent to the homework set’s upload date.

Have a go at the homework set. Come back here next week to check uploaded solutions! <3

HSC Preliminary Chemistry Syllabus Notes

Properties and Structure of Matter

Week 2 Notes – Atomic Structure & Atomic Mass

*Your ConquerHSC Notes are constantly being revised throughout 2019 to ensure quality.

Learning Objective #1 – Investigate the basic structure of stable and unstable isotopes 

  • Their position in the periodic table
  • The distribution of electrons, protons and neutrons in atom
  • Representation of the symbol, atomic number and mass number (nucleon number)

Learning Objective #2 – Model the atom’s discrete energy levels, including electronic configuration and spdf notation 

Learning Objective #3 – Calculate the relative atomic mass from isotopic composition

Learning Objective #4 – Investigate energy levels in atoms and ions 

Collecting primary data from a flame test using different ionic solutions of metals

Examining spectral evidence for the Bohr model and introducing the Schrodinger model

Learning Objective #5 – Investigate the properties of unstable isotopes using natural and human-made radioisotopes as examples of:

  • Types of radiation
  • Types of balanced nuclear reactions
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