GCSE Chemistry Notes PDF

Chemical Equations Equations Chemical Equations Chemical reactions can be represented by word equations or equations using symbols and formulae. Key Aims...

Chemical Equations Equations Chemical Equations Chemical reactions can be represented by word equations or equations using symbols and formulae. Key Aims 1. Chemical Equations. Formation of compounds can be represented in a chemical formula. 2. Balancing Equations. The chemical formula shows you the elements involved and the 3. State Symbols. number of atoms involved. An equation has reactants and products. Reactants are substances that react together in a chemical reaction to form the products. Reactants are found before the arrow in a chemical formula and products are found after the arrow. All chemical equations need to be balanced. On either side of the reaction the number of atoms should be the same and this can be done by putting numbers in front of the atoms or compound. Study Mind Tip Balancing Equations Never change the chemical formula to balance the equation. Let’s work through an example to balance an equation: Always add numbers in front! Remember to balance equations logically. Count the 1. Magnesium + Oxygen → Magnesium Oxide. First write the number of atoms of each chemical formula for the word equation. element on each side of the equation and determine which needs balancing. Start off by Mg + O2 → MgO balancing one element, then proceed to balance the next. Don’t try to balance all the 2. First balance the oxygen. Find an element that does not balance elements at once as you will get confused! and add a number in front of on the other side. Mg + O2 → 2 MgO  3. Then balance the magnesium. This is fixing the discrepancy caused by initially balancing the oxygen. Add another number on the other side to finish balancing the equation. 2 Mg + O2 → 2 MgO State Symbols You can also add state symbols to any reactant or product in the chemical equation. State of Matter State Symbol Solid (s) Study Mind Tip Aqueous just means dissolved in water. This is useful when looking Liquid (l) at a chemical equation and determining if there are any solutions. An example could be an aqueous solution of copper Gas (g) (II) sulfate, Aqueous (aq) Below is a reaction that involves all the state symbols: 2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)  Common Molecules & Ions Recognising Common Ions There are several common ions you will come across throughout GCSE Chemistry, which you need to be familiar with. The following tables show the positive and negative ions you should know. Positive Ions Study Mind Tip Ion Charge It will help to learn which metals are in which group. This make it easier to remember the charge Group 1 metals: Li, Na, K, Rb, 1 of the ions. However, there’s Cs, Fr always a periodic table in the exam to refer back to! Group 2 metals: Be, Mg, Ca, Sr, 2 Ba, Ra Group 3 metals: B, Al, Ga, In, Ti 3 Hydrogen (H) 1 Silver (Ag) 1 Study Mind Tip Copper (Cu) 1 Writing the charge of ions is very important when writing out Ammonium (NH4) 1 chemical equations. Iron II (Fe II) 2 Lead (Pb) 2 Zinc (Zn) 2 Iron III (Fe III) 3  Negative Ions Ion Charge Group 5 non-metals: N, P, As -3 Group 6 non-metals: O, S, Se, -2 ? Knowledge Recall Te 1. What is the charge of an Group 7 non-metals: F, Cl, Br, I, ammonium ion? -1 2. What is the charge of a At silver ion? 3. What is the charge of a hydroxide ion? Hydroxide (OH) -1 Nitrate (NO3) -1 Carbonate (CO3) -2 Sulfate (SO4) -2  Elements & Compounds Atoms All substances are made out of atoms. An atom is the smallest unit of all Key Aims matter. An atom is the smallest part of an element that can exist. 1. Atoms. 2. Elements. A molecule is made of a fixed number of atoms, which are covalently 3. Compounds. bonded together. Elements Periodic Table An element is made of one type of atom. Elements are made up of atoms, and each element only has one type of atom. For example, oxygen is only made up of oxygen atoms. Elements are arranged in the periodic table. There are about 100 different elements, which are arranged in the periodic table based on Study Mind Tip their properties. Make sure you are familiar with the differences between Elements are arranged by atomic number. All elements are arranged elements and compounds as the whole of Chemistry GCSE is in the periodic table in ascending order of atomic number. Elements built on these concepts. are represented using a symbol which is usually one or two letters for e.g. C = Carbon and Na = Sodium. Compounds What is a compound? A compound is a substance that contains two or more elements chemically combined together. The elements in each compound are present in fixed proportions - for example in carbon dioxide (CO2) for every gram of carbon there is 2g of oxygen. The formation of compounds involves breaking and making bonds. Study Mind Tip Making a compound involves making bonds between atoms. We will learn about electron Sometimes you also need to break bonds in the reactants. usually transfer and making/breaking bonds in a lot more detail later results in a change in energy which can be detected. So either energy on. is required for the formation or energy is made during the formation.  Making and breaking bonds involves electrons. Electrons can either be shared, lost or gained to form chemical bonds. For example, covalent bonds are formed when electrons are shared, and ionic bonds are formed when electrons are transferred from metal to non-metal.  Atomic Models Part 1 Models for Atoms Models used to explain atoms have changed overtime as scientists gather new evidence through conducting experiments. We will first look Key Aims at a summary of the first two stages in developing the model for an atom accepted today. 1. Billiard Ball Model 2. Plum Pudding Model Billiard Ball Model (John Dalton) In the early 1800s, John Dalton conducted many experiments, making the following findings to form the Billiard Ball model: Elements are made of tiny particles called atoms. Atoms were spherical in shape and could not be divided or split. Study Mind Tip During chemical reactions, atoms can be combined, separated or rearranged. Try to learn each stage of the progression of the atomic model systematically, so you don’t get Atoms were solid spheres. confused between the different models. Plum Pudding Model (Thompson)  In the late 1800s, Thompson discovered negatively charged particles within the atom, known as electrons. Here is a summary of his findings: Negatively charged electrons are present in atoms. He proposed the ‘plum pudding’ model, which suggested the atom Study Mind Tip was a positively charged sphere with negatively charged electrons randomly embedded within it. Questions about the history of the atom usually focus on the plum pudding model. You need The mass of the atom is equally distributed throughout the atom. to be able to describe why the atom was likened to a plum pudding in detail. Thompson believed that the atom (pudding) was a ball of positive charge, with electrons (plums) embedded within it.  Atomic Models Part 2 Models for Atoms Nuclear Model (Rutherford) Key Aims 1. Nuclear Model. 2. Developed Nuclear Model. Ten years later, Rutherford proposed the nuclear model based on evidence from the alpha particle scattering experiment. Rutherford and his partners (Marsden and Geiger) fired high speed alpha particles, which are densely charged, tiny, positive particles, at a piece of very thin gold foil. Study Mind Tip Rutherford’s particle scattering experiment is one that you need to know. Exam questions tend to focus on the results of this experiment and how it made the model of the atom change drastically. According to the plum pudding model, they were expecting all of the alpha particles to pass through undeflected because the positive charge of the atom was spread throughout the atom.  However, what actually happened was that most alpha particles went straight through the foil, whilst some alpha particles were greatly deflected and scattered. The observations suggested the following conclusions: Observation Conclusion Study Mind Tip Most of the alpha particles went Most of the atom is made up of through the gold foil. empty space. You must understand what we can learn from the observations to draw a conclusion. This will Alpha particles were deflected The nucleus of the atom must be help you to understand how and some rebounded. positive to repel and deflect alpha each atomic model was developed. particles. Only a very small proportion of The nucleus is very small the alpha particles were compared to the size of the atom. deflected. To explain these findings Ernest Rutherford proposed the nuclear model, in which there are electrons that orbit the central positively charged nucleus. Study Mind Tip In the nuclear model here is a tiny positive nucleus, surrounded by Whenever we look at empty space. The electrons are a long way from the nucleus in an outer contrasting scientific theories, try to make a table of the orbit. Most of the mass of the atom is in the central nucleus. differences between the two like this one. This will help you in 6 mark questions that ask you to If Thomson’s plum pudding model was correct, then the alpha particles ‘Compare and contrast…’. should have just passed through. But in fact most alpha particles passed through, but some were deflected greatly by the positive charge in the  1:1 GCSE CHEMISTRY TUTORING Delivered by our expert GCSE Chemistry tutors, who achieved the top grades in the exam themselves. A personalised 1:1 approach, tailored to your unique needs Lessons supported by comprehensive topic-by-topic tutorial slides, past papers, and revision notes Book your FREE consultation now https://studymind.co.uk/subject /gcse-chemistry-tutors/  centre of the atom. There is just empty space between the nucleus and the electrons. Study Mind Tip It is important to remember that these are old models and not what we actually use in theory today. Positive charge is found in the centre (red) and electrons are found around (green). Developed Nuclear Model (Bohr) Neils Bohr adapted Rutherford’s nuclear model by suggesting the Study Mind Tip electrons are at specific distances from the nucleus organised in energy levels. The Bohr model is the one we use at GCSE. You will soon learn more about electron shells and Bohr’s theoretical calculations supported the experimental observations. how we show the number of He suggested there are two electrons in the first shell, up to eight electrons in each shell. electrons in the second shell, and up to eight electrons in the third shell. Bohr suggested that electrons orbit at specific distances (shells) from the nucleus.  Discovery of Protons & Neutrons Key Aims 1. Discovery of Protons. Protons and Neutrons 2. Discovery of Neutrons. Discovery of Protons At this stage, the model showed that there was a centre of positive Study Mind Tip charge, with electrons orbiting in different levels. However, the centre of Make sure to have a thorough positive charge was not understood fully. understanding of protons, including charge, mass and position in the atom. Exam Further experiments led to show that the nucleus could be subdivided questions will have lots of into smaller particles, each having the same positive charge. These questions relating to these aspects. individual positively charged particles were called protons. Discovery of Neutrons Study Mind Tip In 1932 James Chadwick found evidence to support the existence of As with protons, you need to know a neutron’s charge, mass neutrons. He suggested that the nucleus consists of neutrons which and position in the atom. The have a mass but no charge. This was roughly 20 years after the concept questions about neutrons will often come with questions of a nucleus was accepted within the scientific community. about protons.  Study Mind Tip Exams can include a 6 mark question on this topic. They are testing your understanding of how the model of the atom changed over time, so try to order your answer in a logical fashion. This model showed that an atom consisted of electrons orbiting around a central nucleus, consisting of protons and neutrons. Neutrons have no charge but have a mass.  Developing the Atomic Model Looking Forward Overtime scientists have gathered new experimental data which leads to Study Mind Tip the development or replacement of the old models used to explain atoms. For instance after the alpha particle scattering experiment there You need to learn each of these models off by heart and explain was a change in the atomic model from the ‘plum pudding’ model to the them. You need to explain the nuclear model. experimental evidence which led to their development where applicable. Learn Rutherford’s alpha particle scattering experiment off by heart and Worked example: Compare and contrast the plum pudding model with how the observations relate to the nuclear model of the atom. (3 marks) the given conclusions. 1. Compare the structure of each model. In the plum pudding model, there is a positive sphere with negative charge randomly placed within the sphere. There is no empty space. Study Mind Tip In the nuclear model, there is a central positive, tiny nucleus, with mostly Remember for 3 marks you empty space and lots of negative charges a long way from the nucleus. need to give 3 clear points. Try not to waffle and write too much unnecessary information. 2. Compare the masses. The mass in the plum pudding model is distributed equally throughout the sphere. The mass in the nuclear model is concentrated in the nucleus.  Representing Elements Representing Elements in the Periodic Table In the periodic table, each element is represented using a symbol and is surrounded by two numbers. The number at the top of the symbol is the mass number, which is the number of protons plus the number of neutrons within the atom. Recap The number at the bottom of the symbol is the atomic number, Elements are made up of which is the number of protons within the atom. atoms, and each element only has one type of atom. We learnt that there are protons, neutrons and electrons which make up an atom. Now we will look at how these can be represented using atomic symbols. Figure X. Argon. The element argon is represented as the symbol Ar within the periodic table. The number at the top represents the mass number (40) and the number at the bottom represents the atomic number (18). Practice Question:  1. Be2+ has 7 electrons 2. H+ has no electrons 3. Se2- and Fe together have 77 electrons. 4. The atomic number of Be2+ is different to that of Be. 5. The mass number of H+ is different to that of H. Which of the above statements are true? 1. False - Be2+ has 2 electrons. As an uncharged element, it has 4 electrons. However, as it is positively charged with a +2 charge, it loses 2 electrons (4-2=2) 2. True - H+ has no electrons. This is correct, as H only has 1 electron. When it’s positively charged, it loses this electron. Study Mind Tip 3. False - Se2- and Fe together have 62 electrons. Se2- has 34 + 2 = It’s really important to grasp the concept of charges. A positive charge indicates an element or 36 (as 2 electrons gained due to -2 charge). 36 + 26 (electrons in compound has lost electrons. On the other hand, a negative Fe) = 62. charge indicates an element or compound has gained electrons. 4. False - The atomic number of Be2+ is the same as that of Be, as the atomic number is the number of protons, which do not change with electron number changes. If the proton number changes, the element would differ. 5. False - The mass number of H+ is the same as that of H, as there are the same number of protons and neutrons in both uncharged and charged forms of H.  Isotopes Isotopes What are isotopes? Isotopes are atoms of the same element with the Key Aims same number of protons but different number of neutrons in the nucleus. So isotopes have the same atomic number but different mass 1. Isotopes. 2. Examples of Isotopes. numbers. Physical properties of isotopes are different. Isotopes can have varying physical properties, because mass determines physical properties such as density, boiling and melting point. Chemical properties of isotopes are pretty similar. Chemical properties are determined by the number of electrons, which is the same in isotopes. Therefore they have the same chemical properties. Study Mind Tip Isotopes can be radioactive or non-radioactive. Radioactive isotopes There may be an exam have many uses. Common uses include use in medical imaging or in question asking you to describe how the physical and chemical carbon dating. properties of isotopes differ. Example: Isotopes of Chlorine Chlorine-35 and Chlorine-37 have the same number of protons (and electrons), but have a different number of neutrons. Therefore the mass number is different. They are both the same element, so have the same atomic number. Atomic Mass Protons Electrons Neutrons Number Number Study Mind Tip Remember you can work out Cl-35 17 35 17 17 18 the number of neutrons by the mass number minus the atomic number. Cl-37 17 37 17 17 20  Example: Isotopes of Hydrogen Study Mind Tip Isotopes of hydrogen are asked about very often. It may seem a bit strange that Protium has no Atomic Mass electrons, but don’t worry it is Protons Electrons Neutrons correct! Number Number Protium 1 1 1 0 1 Deuterium 1 2 1 1 1 Tritium 1 3 1 2 1 Example: Isotopes of Carbon Atomic Mass Protons Electrons Neutrons Number Number Carbon-12 6 12 6 6 6 Carbon-13 6 13 6 7 6 Practice Question: Magnesium has several different isotopes, including Mg-24, Mg-25 and Mg-26. Which of the following statements about Magnesium’s isotopes are true: Study Mind Tip 1. They have the same number of protons The best way to work through this question is to draw a table 2. They have the same number of neutrons like the ones shown above and 3. They have a different number of electrons write out the numbers. Then it is much easier to work through 4. They have a different mass the statements and determine 5. They have the same atomic number which are true. 6. They have similar chemical properties. 1. Isotopes only differ in neutron number, not protons. 2. Different isotopes have different numbers of neutrons 3. They have the same number of electrons, as their atomic numbers are the same. 4. This is as their mass number differs, due to the variation in neutron number. 5. They have the same atomic number, as the number of protons is invariant. 6. They have similar chemical properties, as the electron number is the same.  Size & Mass of Atoms Size of an Atom Atoms are extremely small. Each individual atom is incredibly small Key Aims and can not be seen by the human eye. The radius of each atom is around 0.1nm, which in standard form is 1 x 10-10 m. 1. Size of an Atom. 2. Mass Number. 3. Calculations of Protons, The nucleus is only a tiny part of an atom. Each atom has a central Neutrons & Electrons. nucleus and the radius of this nucleus is around 1 x 10-14 m. The nucleus is tiny compared to the whole atom. The radius of the nucleus is less than 1/10,000 of the radius of the whole atom. This means that most of the atom is occupied by the cloud of electrons. Study Mind Tip You need to learn the radius of the atom and nucleus off by heart for your exams.  1:1 GCSE CHEMISTRY TUTORING Delivered by our expert GCSE Chemistry tutors, who achieved the top grades in the exam themselves. A personalised 1:1 approach, tailored to your unique needs Lessons supported by comprehensive topic-by-topic tutorial slides, past papers, and revision notes Book your FREE consultation now https://studymind.co.uk/subject /gcse-chemistry-tutors/  Fig X. Size of an atom. The central nucleus (red) is extremely small compared to the size of Rutherford’s nuclear model. Positive charge is found in the centre (red) and electrons are found around (blue). (http:// dispatchesfromturtleisland.blogspot.com/2013/04/an-atom-drawn-to- scale.html) Mass Number The mass number is the total number of protons and neutrons which are found in the nucleus. Electrons have a negligible mass, so we don’t usually count it when working out the mass of an atom. Study Mind Tip Protons, neutrons, electrons have a relative mass. Both protons and Electrons actually has a mass of neutrons have a relative mass of 1. Electrons have a really small mass around 1/1800, but this is such so it is usually assumed as 0. a small number we can just assume it is zero when doing calculations. Proton Neutron Electron Mass 1 1 Very small The majority of the mass of an atom is in the nucleus. The nucleus consists of the majority of the mass of an atom because the subatomic particles within the nucleus (the protons and the neutrons) both have a relative mass of 1. Whereas the electrons surrounding the nucleus have a very small mass; the mass is negligible. Study Mind Tip Calculations of Protons, Neutrons, Electrons Becoming familiar with We can perform calculations to determine the number of protons, calculating the numbers of protons, neutrons and electrons neutrons and electrons within an atom: in an atom or ion, given its atomic number and mass number. Worked Example: Use your periodic table to calculate the number of protons, neutrons and electrons in phosphorus. Answer: 1. Find phosphorus on the periodic table. It has a mass number of 31 and an atomic number of 15 2. Number of protons = number of electrons. Protons = 15, therefore electrons = 15 3. Number of neutrons = Atomic mass number - Atomic number 31-15= 16 neutrons  Relative Electrical Charges Atoms and Elements The nucleus contains protons (+1 charge) and neutrons (0 charge). Electrons (-1 charge) orbit in shells around the central nucleus. The mass if concentrated in the nucleus. Recap We learnt that elements are The number of protons = the number of electrons in a non-charged made of lots of small atoms. We atom. also learnt more about the history of research into atomic structure. By the end of the tutorial, we learnt about the Protons, Neutrons, Electrons current known structure of the atom, summarised here. Structure of an Atom An atom consists of neutrons, protons and electrons. Each atom has 2 regions: The nucleus is found in the centre of the atom, and consists of positively charged protons and uncharged neutrons. This makes the atomic nucleus positively charged. A cloud of electrons surrounding the nucleus. The electrons are negatively charged and orbit the nucleus in shells. Study Mind Tip Relative Charge Make sure you learn the relative charges of protons, neutrons Protons, neutrons and electrons have relative charge. We can give a and electrons. relative charge to protons (+1), neutrons (0) and electrons (-1). Make sure to memorise these! All atoms have no overall charge. Atoms are neutral and have no overall charge because the number of protons (positively charged)  equals the number of electrons (negatively charged). The opposite charges cancel each other out. Proton Neutron Electron Relative Charge 1 0 -1 Study Mind Tip Fig 1. Relative Charges. Each of proton, neutrons and electrons have a relative charge. If two atoms don’t have the same number of protons, then they are not from the same element. We will later learn about isotopes, which are Atomic Number different atoms of the same element which have the same number of protons but different The atomic number is the number of protons in the nucleus. Each number of neutrons. So the element has its own atomic number and specific characteristic that helps number of neutrons can vary between atoms of the same you recognise it. The periodic table is arranged in ascending atomic element, but never the number number. of protons. Atoms of the same element have the same number of protons. The number of protons in a particular atom is equal to the atomic number and all atoms of a specific element have the same number of protons. For example, calcium has an atomic number of 20. Therefore calcium has 20 protons and all atoms of calcium have the same number of protons - 20. Atoms of different elements have different numbers of protons. The atomic number of each element is different, therefore the number of protons each element has is different. For example, calcium has an atomic number of 20 and therefore has 20 protons. However sodium has an atomic number of 11 and therefore has 11 protons.  Relative Atomic Mass Understanding RAM The relative atomic mass (Ar) is the average mass of atoms of an element relative to the mass of an atom of carbon-12 (which is given a mass exactly of 12). The average mass must take into account the Key Aims proportions of naturally occurring isotopes of the element. 1. Understanding RAM. 2. Calculating RAM. Scientists decided that the atomic mass of an atom of an element would be compared to carbon-12. Therefore the atomic mass is known as relative as it is being compared to the mass of carbon-12. If the Ar of an atom is lower than 12 it has a mass smaller than carbon-12 atom. You can find the relative atomic mass of an element on a periodic table by looking at the number directly above the element symbol. For example the relative atomic mass of Copper (Cu) is 29. Calculating RAM Study Mind Tip The relative atomic mass of an element is the average mass of an atom, You should be able to calculate and it takes into account the masses of each isotope and their the relative atomic mass of an proportions in the environment. element given the percentage abundance of its isotopes. To calculate the relative atomic mass you require the following information: The abundance of each isotope which tells you the quantity the isotope is present in the environment. This is different for each isotope of an element. The mass number of each isotope of a particular element. The number of protons in all the isotopes of an atom remains the same, it is only the number of neutrons. Lets work through an example:  Practice Question: A sample of chlorine gas is a mixture of 2 isotopes, chlorine-35 and chlorine-37. These isotopes occur in specific proportions in the sample i.e. 75% chlorine-35 and 25% chlorine-37. Calculate the relative atomic mass of chlorine in the sample. 1. Identify the isotopes. Both isotopes present in the mixture or environment, and the abundance of each isotope of the element. Relative isotopic Abundance Cl-35 75% Study Mind Tip Cl-37 25% Relative atomic mass is may not be written as a whole number due to the fact that is often an average of the masses of many isotopes. 2. Identify the mass number. For each isotope of the element. Relative isotopic Mass Number Abundance Cl-35 35 75% Cl-37 37 25% 3. Write out the formula to calculate relative atomic mass. 4. Substitute identified data. Substitute into the formula and work out the answer. = 35.5  The relative atomic mass of chlorine in this mixture was 35.5. This number is closer to 35 compared to 37 as chlorine-35 is more abundant than chlorine-37. We can do a sense-check, to make sure that the value seems right. 35.5 is in between 35 and 37, as we would expect, and it has a closer mass to Cl-35, which is the more abundant isotope. This seems fine! You need to be able to rearrange the equation to work out a missing abundance or mass. Study Mind Tip This type of question comes up Worked example. Oxygen has three isotopes. The abundances in frequently. Make sure you learn percentage are given here. The mass of one of the isotopes is unknown. the equation and know how to rearrange it. It is important that The average atomic mass of the isotopes is 16.65. Work out the values of you are confident with using x and y. algebra and rearranging equations to find an unknown value. Answer: 1. Set up a table. Put in all the values as shown.  Mass x Isoptope Mass Abundance (%) Abundance 16 O 16 50 16 x 50 = 800 17 O 17 35 17 x 35 = 595 x = 100 - yO y y x 15 (50+35) = 15 2. Work out x. All percentages = 100, so x = 100-(50+35) =15% 3. Form an equation for y. Use the given atomic mass and the equation for calculating it. 16.65 = (800 + 595 + 15y) / 100 4. Rearrange to find y. [16.65 x 100 - (800+595)] / 15 = y = 18 So the atomic mass of the unknown isotope = 18 and the abundance = 15%  Developing the Periodic Table The Early Periodic Table By the 19th century scientists had discovered over 50 elements, many attempts had been made to put these elements into a logical order to reflect the similarities in their chemical properties. Key Aims 1. The Early Periodic Table. The early periodic tables were arranged strictly by atomic weight as 2. Newlands Octaves. protons, neutrons and electrons had not been discovered yet. Below is a 3. Mendeleev’s Periodic Table. summary of how the periodic table has changed over time to form the periodic table we use today. Newlands Octaves Newton’s Law of Octaves. In 1864, an English scientist, John Newland proposed his ‘law of octaves’ where he arranged the elements in order of their atomic mass. Newland noticed that every 8th element had similar chemical properties. Study Mind Tip Newland’s Octaves. Newland arranged the discovered elements in You need to be able to interpret order of atomic weight with H (Hydrogen) as the lowest then Li and compare the ideas of early (Lithium) then Be (beryllium) and so on. Newland proposed that scientists in developing the periodic table. Look carefully at properties of every 8th element were similar for example Li (lithium) the types of elements that have and Na (sodium). been grouped together. Problems with the Law of Octaves. However, many scientists did not accept Newland’s law of octaves because many new elements were still being discovered and did not fit in the table. Furthermore, all elements in his octaves did not have similar properties for example in the 7th octave O (Oxygen a non-metal) and Fe (Iron a metal). Mendeleev’s Periodic Table  In 1869 Dimitri Mendeleev categorised the elements into ‘Periodic System’. This was a table of elements arranged according to atomic mass however there were many differences between Newland’s octaves and the Mendeleev’s periodic system. Mendeleev’s main focus was arranging the elements based on similarities Study Mind Tip of chemical and physical properties. So he arranged the elements in a table, with vertical columns known as groups and these groups had Learn the differences between Newland’s octaves and elements with similar properties. Mendeleev’s periodic system. Mendeleev acknowledged that all elements had not been discovered yet, therefore left gaps to be filled in once they had been discovered. The gaps existed in certain groups with particular characteristics, therefore Mendeleev was able to predict the properties of these undiscovered elements. When these elements were discovered, his predictions were right and Mendeleev’s table was accepted by other scientists. Rearranging the Position of Elements He did not stick to the strict order of arranging elements in order of increasing atomic mass. When Mendeleev organised elements in order of atomic mass, some elements did not fit the pattern of the group of increasing atomic mass. He rearranged these elements so Study Mind Tip that they would be placed in a group with similar properties. When it comes to looking at the periodic table, always think Iodine and Tellurium swapped positions. Iodine has a lower atomic about the different physical and chemical properties of the mass than Tellurium, therefore it should be placed before Tellurium in elements in each group. the periodic table. However, Mendeleev saw that Iodine has similar properties to the elements in Tellurium’s group (7) such as fluorine, chlorine and bromine. Therefore, Mendeleev swapped around these two elements to ensure elements were in groups with similar properties. Explanation of Tellurium and Iodine order. We now know that the explanation of isotopes is the reason why iodine has a lower mass number than tellurium, even though it has a higher proton number.  The Current Periodic Table The Modern Periodic Table Discovery of isotopes of the same element explained the order based on atomic mass was not always correct. When scientists Key Aims discovered isotopes it explained why some atoms had heavier atomic 1. Modern Periodic Table. masses than expected. This meant that Mendeleev was correct to not 2. Arrangement of the Periodic Table. stick to arranging the periodic table according to atomic mass. Elements in the modern periodic table are arranged according to atomic number. In the 20th century, protons and electrons were discovered by scientists who then realised the periodic table should be arranged according to atomic number rather than atomic mass. Order of elements in modern periodic table. Elements in the modern periodic table are arranged according to atomic number. The periodic table we use today is arranged according to the number of protons an element has also known as the atomic number. Before the 20th century, scientists had not discovered subatomic particles such as protons and electrons. Therefore scientists arranged elements in order of atomic weight in the past before this discovery was made. Study Mind Tip The table is called a periodic Arrangement of the Periodic Table table because similar properties occur at regular intervals. Elements with similar properties are arranged in vertical columns known as groups. In the periodic table, elements are arranged in vertical columns’; these are known as groups. Each group of elements has similar chemical and physical characteristics. Let's take group 1 all the elements. in group 1 such as lithium, sodium and potassium are all very reactive alkali metals.  Groups = column down Study Mind Tip 1 2 3 4 5 6 7 You should be able to predict possible reactions and probable reactivity of elements from their H Li Be B C N O positions in the periodic table. With more practice you will F Na Mg Al Si P S become more familiar with the similarities and differences across the elements in the Cl K Ca Cr Ti Mn Fe periodic table. Elements in the same group have the same number of electrons in their outermost shell. Elements in each group in the periodic table have a distinctive set of properties and react in the same way, this is due to the number of electrons in the outermost shell. When atoms of elements react, they either gain, lose or share the electrons in their outermost shell. Elements in the periodic table are also arranged in horizontal rows. The periodic table is formed of multiple rows, these are known as periods. The elements are arranged in periods in ascending atomic numbers. Each period signifies the number of shells that the atoms of Study Mind Tip the elements in that period have. For example, lithium is in the second period and has 2 shells. You need to be able to “read” the periodic table and determine the number of Group Number. The group number at the top signifies the number of electrons in the outer shell and the number of electron shells electrons in the outermost shell. from the group number and period number from the Periodic Table. Period.There are also horizontal rows known as periods. These rows tell you the number of electron shells that an atom has. Fig 1. Periodic Table. The modern periodic table is arranged by ascending atomic number. It is made of vertical columns known as groups.  Electronic Configuration Electronic Configuration Working out the Electronic Configuration As we mentioned in a previous tutorial, a cloud of electrons surround the nucleus. These electrons moved in different energy levels, which we can call electron shells. We can work out the arrangement of electrons into shells. This is called Study Mind Tip the electronic configuration. The first shell will always have two electrons. Electrons occupy the closest shell first. Electrons will occupy the shell closest to the nucleus before occupying the next shell. The shell must be full before another shell is occupied. The shells hold different numbers of electrons. The first shell holds up to two electrons. The second and third shells can hold up to eight electrons. The electronic configuration can be written. We can write the electronic configuration by writing the number of electrons in each shell, separating the numbers with commas. For example electronic configuration of potassium is 2, 8, 8, 1. Study Mind Tip When drawing the electron shells, the electrons are represented as crosses. It is best to draw them at equal distances from each other on each shell. Fig 1. Potassium. Electron shells of potassium. Potassium has an electronic configuration of 2, 8, 8, 1. Relating the Electronic Configuration to the Periodic Table The number of notations determine the period. For example the electronic configuration of potassium is 2, 8, 8, 1. There are four notations, therefore potassium is in period 4. The number of the last notation determines the group. The last notation for the electronic configuration of potassium is 1. Therefore potassium is in group 1.  Metals & Non-Metals Formation of Ions Ions are formed from atoms when electrons are transferred between atoms: Key Aims A positive ion is formed when an atom has lost electrons, after which 1. Formation of Ions. 2. Properties of Metals and it contains less electrons than protons. Atoms of elements that react Non-Metals. to form a positive ion are metals. A negative ion is when an atom gains electrons so it contains more electrons than protons. Atoms of elements that react to form a negative ion are non metals. Let's take sodium and chlorine as an example. When sodium reacts with chlorine, it losses the outermost electron and becomes a positively charged sodium ion. This suggests that sodium is a metal. Study Mind Tip If you get confused when working out the charge of an ion, draw out some electron configuration diagrams. This will help you to work out how many electrons will be lost or gained. Fig 1. Sodium ion formation. Sodium atoms have one electron in their outermost shell with an electron configuration of 2,8,1.  When the sodium atom loses the electron in the outermost shell it becomes a sodium ion. The ion now has 2 full shells and has a positive charge. The electronic configuration is now [2,8]+ Study Mind Tip When Chlorine gains the electron lost by sodium and becomes a negatively charged chloride Ion. This suggests that chlorine is a non Remember electrons are negatively charged. So if an metal. electron is gained, the ion will have a negative charge. Fig 2. Chloride ion formation. Chlorine atoms have 7 electrons in their outermost shell with an electron configuration of 2,8,7. When the chlorine atom gains the electron in the outermost shell it becomes a chloride ion. The ion now has 3 full shells with a negative charge. The electronic configuration is now [2,8,8]- Study Mind Tip It is really important that you Properties of Metals and Non-Metals are able to explain the differences between metals and Metals and non-metals have different physical and chemical properties. non-metals on the basis of their characteristic physical and Below is a table that summarises these main differences chemical properties. A lot of these properties are covered in Physics GCSE as well, particularly the conduction of heat and electricity.  Metals Non-metals Appearance Shiny Dull Strength Strong Weak Density High Low Solid Solids and Gases State at Room exception: mercury - exception: bromine - Temperature liquid liquid Melting and High Low Boiling Points ? Knowledge Recall Conduction of Good conductors of Poor conductors of heat 1. Define a positive ion. Heat heat good insulators 2. Define a negative ion. 3. Are metals good or bad Conduction of conductors of heat? Good Poor 4. Do metals generally have a Electricity high or low density? Ductile (stretched Ductile Non-ductile into wires) Malleable (beaten Malleable Brittle into sheets) Metals as Conductors Metals are good conductors of both heat and electricity because they contain delocalised electrons which are free to move around. Delocalised electrons can carry current, throughout the structure of the metal, as well as transfer kinetic energy between themselves.  Periodic Table Metals vs Non-Metals Metals and Non-Metals Key Aims The metal and non-metals in the periodic table can be further 1. Group 1 and 2 Metals. 2. Transition Metals. subdivided. Below we will discuss these divisions and the physical and 3. Group 0. chemical properties that group them. Group 1 and 2 metals Metals in groups 1 and 2 are known as the reactive metals. During a reaction, elements in Group 1 lose 1 electron and all react in the same way. All elements in Group 2 elements lose 2 electrons and react in the same way. Study Mind Tip We will learn more about the reactions that Group 1 and 2 elements undergo soon. Here is a brief overview of the reactive metals’ properties: You should be able to explain how the atomic structure of metals and non-metals relates They react vigorously with other elements such as oxygen and to their position in the periodic table A common exam question chlorine. is having to explain how the reactions of elements are related to the arrangement of They react with water to form alkaline solutions and are also known as electrons in their atoms and alkaline metals. hence to their atomic number. They are soft and can be cut very easily with a knife. Transition Metals Transition metals are the blocks of elements found between groups 2 and 3, in the centre of the periodic table. These are the elements that most people associate the world metal to such as copper, iron and silver. They are very strong, good conductors of electricity and heat. Group 0  Non-metals in Group 0 are also known as noble gases. These elements do not react easily and this is because they have a full outer shell of electrons. We will learn more about transition metals and their properties in a later tutorial. Study Mind Tip Make sure you are familiar of the positions of the different groups and the split between metals and non-metals. Fig 1. Subdivisions in the Periodic Table. The group 1 and 2 reactive metals (red and orange) are on the left-hand side of the periodic table. The transition metals (pink) are between the main grouped elements and are at the centre of the periodic table. The noble gasses are furthest to the right-hand side of the table (blue).  1:1 GCSE CHEMISTRY TUTORING Delivered by our expert GCSE Chemistry tutors, who achieved the top grades in the exam themselves. A personalised 1:1 approach, tailored to your unique needs Lessons supported by comprehensive topic-by-topic tutorial slides, past papers, and revision notes Book your FREE consultation now https://studymind.co.uk/subject /gcse-chemistry-tutors/  Group 1 Group 1 - The Alkali Metals Group 1 are metals that are found on the left hand-side of the periodic Key Aims table. They are also known as the alkali metals and are highly reactive. 1. Group 1. Through this tutorial we will be exploring the physical and chemical 2. Physical Properties. properties of the alkali metals and the trends through group 1. 3. Chemical Properties. Physical properties of the alkali metals The alkali metals are soft solids at room temperature and can easily be cut with a knife. They also have low densities which means that these elements can float on water. They have low melting and boiling points; this decreases as you go Study Mind Tip down the group. There may be an exam question asking you to explain Chemical properties of the alkali metals the different physical and chemical properties of the group 1 alkali metals. Atoms of group 1 elements react in the same way. All the elements in group 1 are highly reactive, have the same properties and react in the same way this is because all the alkali metals have a single electron in their outermost shell. When atoms of group 1 elements react, they each lose one electron forming a positively charged ion. Fig 1. Group 1 Ions. When atoms of group 1 elements react they lose 1 electron. For example, when lithium reacts it loses 1 electron forming a +1 lithium ion. All elements in group 1 react in the same way and have similar properties.  Group 1: Reactions Group 1 - reactions with oxygen All group 1 metals react quickly with oxygen in the air to produce a metal oxide. Below is a generic formula that represents the alkali metals Key Aims (X) reacting with oxygen to form a metal oxide: 1. Reactions with Oxygen. 2. Reactions with Water. 4 X (s) + O2 (g) → 2 X2O (s) 3. Reactions with Chlorine. Group 1 elements are stored in oil this is to prevent them from reacting spontaneously with oxygen. When these metals are left out, the oxygen in the air reacts with the surface of the metal forming a white oxide that coats the surface. They can also be heated in glass jar full of oxygen, where they burn strongly with a flame forming a metal oxide. Some elements form mixtures of a metal oxide (X2O), metal peroxide (X2O2) and metal Study Mind Tip superoxide (XO2). Below is a table describing the observations made The reactions of the first three when the first three elements in group 1 react with oxygens and the alkali metals with oxygen, equations. chlorine and water should become second nature to you after lots of practice. Observations Equation Red Flame - forms lithium 4 Li (s) + O2 (g) Lithium oxide → 2 Li2O (s) 4 Na (s) + O2 (g) Orange Flame - forms → 2 Na2O (s) Sodium sodium oxide and sodium peroxide 2 Na (s) + O2 (g) → 2 Na2O2 (s) 2 K (s) + O2 (g) Lilac Flame - forms potassium → 2 K2O2 (s) Potassium peroxide and potassium superoxide K (s) + O2 (g) → KO2 (s)  Group 1 - reactions with water All group 1 metals react vigorously with water to produce a metal hydroxide and hydrogen. The Group 1 metals floats on the surfaces and Study Mind Tip fizzes vigorously. Below is a generic formula that represents the alkali metals (X) reacting with water: If you memorise the generic equation and practice balancing it, then it will be easier for you to write the equations for 2 X (s) + 2 H2O (l) → 2 XOH (aq) + H2 (g) reactions with specific elements. These metal hydroxides (XOH) are soluble in water and represented as being aqueous (aq). They form a colourless solution. When a universal indicator is added to measure the pH of the solution it turns purple, this suggests that the solution is highly alkaline. Hence, group 1 elements are also known as alkali metals. Below is a table describing the observations made when the first three elements in group 1 react with water and the equations. Observations Equation Floats on the surface, effervescing 2 Li(s) + 2 H2O (l) → Lithium gently. Gradually become smaller as it reacts and until it disappears. 2 LiOH (aq) + H2 (g) Study Mind Tip Floats on the surface, effervescing Memorise the observations as rapidly. Quickly, melts into a ball 2 Na(s) + 2 H2O (l) → well as the equations. An exam Sodium gets smaller and smaller until it 2 NaOH (aq) + H2 (g) question may tell you the observation and you will have to disappears. determine with metal was used in the experiment. Burns violently giving off a purple 2 K(s) + 2 H2O (l) → Potassium flame. Disappears rapidly, sparks are also seen. 2 KOH (aq) + H2 (g)  Group 1 - reactions with chlorine All group 1 metals react vigorously with chlorine gas forming white salts or colourless crystals known as metal chlorides. Below is a generic formula that represents the alkali metals (X) reacting with chlorine: Study Mind Tip 2 X (s) + Cl2 (g) → 2 XCl (s) Don’t forget to use the state symbols to distinguish between All alkali metals have 1 electron in their outermost shell. During this solids, liquids, gases and reaction the alkali metal X loses the electron in its outermost shell and aqueous solutions. forms a +1 ion. Chlorine then gains that electron, to form a chloride ion (Cl-). The oppositely charged ions are attracted to each other and form a metal chloride (XCl).  Group 1: Reactivity Group 1 - trend in reactivity Trend in reactivity in a group can be explained using the electronic structure of atoms. Looking at the reactions discussed beforehand, when going down group 1 the reactions get more and more vigorous. This suggests that as you go down group 1, the reactivity of the elements increases. So lithium would be the least reactive and francium would be the most reactive out of all the elements in group 1. Study Mind Tip You should be able to explain how properties of the elements Explaining the trend in reactivity in Group 1 depend on the outer shell of electrons of the atoms and predict properties from As you go down the group the reactivity of the Group 1 increases given trends down the group. because: The atoms become larger. This is due to the number of shells increasing. Therefore, the outer electron is further away from the positive nucleus, as you go down the group. The electrostatic attraction of the outer electron to the positive nucleus decreases as you go down group 1. The number of electron shells increases. This increases the shielding effect the positively charged nucleus has on the outer electron. Therefore, the attraction of the outer electron to the nucleus decreases as you go down group 1. Study Mind Tip Therefore, as you go down the group the outermost electron becomes The explanation for the easier to lose and the reactivity increases as you go down the group. reactivity of Group 1 elements down the group is a common six mark exam question. Learn all the factors off by heart.  Study Mind Tip By learning the trend in reactivity, you should be able to predict other reactions, for example with astatine which is the least chemically reactive of the halogens. Looking at the reactions and observations from the previous tutorial, we can predict that astatine would react very slowly with hydrogen and only a small amount of hydrogen astatide would form. Fig 1. Group 1 atoms. As you go down group 1, the atoms become larger and the outer electron is further away from the nucleus. There are also more electrons shielding the positive charge of the nucleus as you go down the group. The force of attraction between the nucleus and the other electron decreases and the electron can be easily lost as you go down the group.  Group 7 Group 7 - The Halogens Elements that are non-metals are found on the right hand-side of the periodic table. They are also known as the halogens. Through this tutorial we will be exploring the physical and chemical properties of the halogens and the trends through group 7. Key Aims 1. Group 7. 2. Physical Properties. Physical properties of the halogens 3. Chemical Properties. How does the mass and boiling point change down Group 7? As you go down group 7, the relative atomic mass increases and the boiling point of halogens also increases. So fluorine has the lowest atomic mass and the lowest boiling point whereas astatine has the highest atomic mass and the highest boiling point. The boiling point is affected by the intermolecular forces. Between atoms of a particular element, there are weak attractive forces known Study Mind Tip as intermolecular forces. When molecules are boiled or melted, these intermolecular forces are overcome. Again, like with the group 1 metals, you should be able to predict properties from given Down group 7, the melting and boiling points increase. The atoms trends down the group. increase in size, as they gain extra electron shells, and the intermolecular forces become stronger. More energy is required to break these forces, thus there are higher melting and boiling points as you go down the group. Halogens have covalent bonding. The halogens are naturally found as simple molecules - a pair of halogen atoms sharing a pair of electrons, this is known as covalent bonding.  Study Mind Tip The reactions across a group will be similar because all the elements will have the same number of electrons in their outer shell. For example, all Fig X. Halogen Molecule. Halogens exist naturally as halogens have seven electrons in their outer shell. a pair of the same halogen atom sharing 2 electrons. In the chlorine molecule, there is a pair of electrons shared between the two atoms. Each chlorine atom donates one electron to be shared. Below is a table that summarises the states of the halogens and their colours. Halogen State Colour Fluorine (F2) Gas Pale Yellow ? Knowledge Recall 1. How does the boiling point Chlorine (Cl2) Gas Green change down group 7? Give an explanation for your Bromine (Br2) Liquid Red-brown answer. 2. What type of bonding do halogens have? Solid 3. What colour is chlorine at Grey Iodine (I2) easily vaporises into a room temperature? gaseous vapour - violet 4. What state of matter is gas bromine at room temperature? Chemical properties of the halogens All atoms of group 7 elements have the same properties and react in the same way, this is because all halogens have 7 electrons in their outermost shell. They need to gain one electron to make a full outermost shell. This can happen in two ways: 1. By reacting with a non-metal and forming covalent compounds and 2. By reacting with a metal and forming ionic compounds.  Group 7: Reactions & Displacement Halogens - Reactions with Non-Metals Halogens also form covalent bonds with other non-metals to form a molecule. Lets work through an example, below is a generic formula that Key Aims represents the halogens (X) reacting with hydrogen: 1. Reactions with Non-Metals. 2. Reactions with Metals. H2 (g) + X2 (g) → 2 HX (g) 3. Displacement Reactions. When this reaction occurs a hydrogen halide is formed which is gaseous at room temperature. When dissolved in water this forms an acidic solution. Halogens - Reactions with Metals The halogens react with metals to form salts. When these salts are dissolved in water they form colourless solutions. Below is a generic formula that represents the halogens (X) reacting with sodium: Study Mind Tip Remember when a halogen 2 Na (s) + X2 (g) → 2 NaX (s) becomes an ion it ends in -ide. Let’s take fluorine, when it gains an electron it becomes an ion and is called fluoride. When a halogen reacts with a metal, the metal atom loses an electron forming a positively charged ion. That lost electron is then taken up by the halogen atom, this forms a negatively charged ion. The oppositely charged ions are attracted to each other; this is known as ionic bonding. Fig 1. Sodium Chloride. Sodium metal loses an electron (blue) to form a positively charged sodium ion.  The electron (blue) is then taken up by chlorine to form a negatively charged chloride ion. 2 Na (s) + Cl2 (g) → 2 NaCl (s) Halogens - Displacement Reaction A displacement reaction is when more reactive halogen can displace a less reactive halogen from a solution of its salt. All these reactions occur in solution and all the salts when dissolved in water are colourless. Let's work through some examples. 1. Chlorine solution reacting with potassium bromide 2 KBr (aq) + Cl2 (aq) → 2 KCl (aq) + Br2 (aq) Chlorine is the more reactive halogen and it will displace bromine from Study Mind Tip potassium bromide. The aqueous bromine will turn the solution from colourless to yellow-orange. Learn these reactions off by heart. Make sure you know the colours produced from each 2. Bromine solution reacting with potassium chloride reaction: bromine forms an orange colour and iodine forms a brown solution. 2 KCl (aq) + Br2 (aq) → NO REACTION Chlorine is the more reactive halogen and will not be displaced by bromine. Therefore no reaction will occur and the solution will remain colourless. Below is a summary of all the displacement reactions you need to know, make sure you learn how to write a balanced equation for all these reactions and observations.  Potassium Potassium Potassium Iodide Study Mind Tip Chloride Bromide Remember that no displacement reactions will 2KCl + Br2 2KCl + I2 occur in some, for example there will be no displacement Chlorine - reaction with chlorine and a Colourless solution Colourless solution chloride solution as their reactivity is the same. to orange to brown 2KBr + I2 Bromine No reaction - Colourless solution to brown Iodine No reaction No reaction -  Group 7: Reactivity Halogens - trend in reactivity As you go down Group 7, the reactivity of halogens decreases so fluorine is the most reactive halogen and astatine is the least reactive halogen. Halogens react to gain an electron so they become more stable. The Study Mind Tip negatively charged electron is attracted to the halogen atom due to the positive charge of the atom’s nucleus. As you go down the group the The way the atoms change down the group are identical to reactivity of the halogens decreases because: group 1. In Group 1 atoms are trying to lose electrons, so they become more reactive down the The atomic mass of the halogens increases. They increase in electron group. In group 7 they are trying shells; so the atoms are larger as you go down the group. Therefore, to gain electrons, so they are less reactive down the group. the attraction of the outer electron to the nucleus decreases as you go Learn the explanation off by down group 7. heart. The number of electrons increase. There are more electrons shielding the positively charged nucleus from the outer electron. Therefore, the attraction of the outer electron to the nucleus decreases as you go down group 7. The weaker electrostatic attraction of the electron being gained to the positive nucleus, makes it harder for the atom to gain an electron and it decreases in reactivity.  Group 0 Group 0 - Noble Gases Group 0 elements are non-metals that are found on the right-hand side of the periodic table. They are also known as noble gases, this is due to Key Aims the highly unreactive nature of these elements. Through this tutorial we will be exploring the physical and chemical properties of the noble 1. Group 0. gasses and the trends through the group. 2. Chemical Properties. 3. Physical Properties. 4. Uses of Noble Gases. Chemical properties of the Noble Gases Noble gases are inert. Group 0 elements are highly unreactive (inert) and do not form molecules readily, this is because the electronic structure of these elements is very stable. The noble Study Mind Tip gasses all have a full outer shell so they do not need to react to Remember that all the noble gain, lose or share electrons. gases have eight electrons in their outer shell, except for helium, which has only two Noble gases exist as single atoms. The noble gases do not form electrons. molecules and are monoatomic. As all noble gasses have a full outer shell of electrons and are very stable, they do not need to share electrons to achieve more stability therefore exist as single atoms. Physical properties of the Noble Gases Boiling point increases down Group 0. As you go down group 0, Study Mind Tip the relative atomic mass increases and the boiling point of noble The noble gasses all have a full gases also increases. So helium has the lowest atomic mass and the outer shell of electrons. All lowest boiling point whereas radon has the highest atomic mass and Group 0 elements, except helium, have 8 electrons in their the highest boiling point. outer shell. However, helium has 2 electrons. Remember before we discussed electronic Explanation for change in boiling point. Between atoms of a structure and the first shell can particular element, there are weak attractive forces known as only hold up to 2 electrons. intermolecular forces. When atoms are boiled or melted, these intermolecular forces are broken. As you go down group 0, the atoms increase in size, as they have more electron shells. The intermolecular forces become stronger so more energy is required  to break these forces thus a higher boiling point as you go down the group. Study Mind Tip In an exam you might be asked to predict the boiling of a particular noble gas given the Uses of Noble Gases boiling point of others. The main use of noble gases is to provide an inert atmosphere. For example argon is used in lamps and light bulbs, as it will not react with the hot filament inside the bulb when it is hot. Also helium is used for filling balloons due to it’s low density compared to air.  The Transition Metals Transition Metals Transition elements are metals that are found in the central block of the periodic table between Group 2 and Group 3. These are the elements that most people associate the world metal to such as copper, iron and silver. During this tutorial we will be discussing the properties of transition metals. When we talk about transition metals, we will be focusing on the properties of chromium, manganese, iron, cobalt, nickel and copper. Study Mind Tip The transition metals listed above have similar properties however are You should be able to exemplify these general properties by different from the alkali metals we discussed before. Below is a table reference to Cr, Mn, Fe, Co, Ni, comparing the properties between Group 1 and transition metals: Cu. Group 1 Metals Transition Metals Melting Point Lower Higher Density Less Dense More Dense Study Mind Tip Strength Weaker Stronger A very common exam question is to describe the difference Hardness Softer Harder compared with Group 1 in melting points, densities, Reacts with oxygen strength, hardness and reactivity Reactivity - Reacts quickly at room with oxygen, water and when heated to form a Oxygen temperature halogens. metal oxide Reacts very slowly or no Reacts vigorously with Reactivity - Water reaction at all with cold cold water water Some transition metals Reactivity - Reacts vigorously with react with halogens but Halogens halogens very slowly  Transition metals have more than one ion. Transition metals form ions with various charges. For example: Copper has 2 ions Cu+ or Cu2+ Study Mind Tip , Iron has 2 ions Fe2+ or Fe3+ and Cobalt has 2 ions Co2+ or Co3+. We would say that transition metals have ‘variable oxidation Transition metals can be used as catalysts. Transition metals can act states.’ This will make more as catalysts which means they increase the rate of reaction without sense when we look at the transfer of electrons. being used in the reaction. For example an iron catalyst is used in the Haber process to produce ammonia.  Separating Mixtures Mixtures and Compounds Mixtures Key Aims A mixture is made from two or more elements or compounds being 1. Mixtures. mixed together, without the formation of any chemical bonds. This 2. Compounds. 3. Separating Mixtures. diagram represents th