2. Atoms, Elements & Compounds (Part 2) (2/4) (Cambridge IGCSE Chemistry 0620 for 2026, 2027 & 2028)

[0:00]Hi everyone, welcome to IGCSE Study Buddy, where you can revise chemistry topics from the Cambridge IGCSE syllabus. If you are enjoying our videos so far, please don't forget to hit the like button and subscribe to our channel. In this video, you are going to learn part two of topic two, Atoms, elements and compounds. Electron configuration is a way of describing how electrons are arranged in an atom. Electrons are arranged in shells around the nucleus. Shells are represented as rings around the nucleus. Each shell can hold a certain number of electrons: The first shell can hold two electrons. The second shell can hold eight electrons. And the third shell can hold eight electrons as well. Electrons fill up these shells or energy levels from the innermost to the outermost. The inner shells fill first before moving to the next shell. So you have to know how to determine the electron configuration of elements and their ions with proton number 1 to 20. So here they are, take a look. Remember, you will be given the periodic table for the exams, so if you want to determine the electron configuration of an element, simply refer to your periodic table and look at the atomic number of the element, which is equal to the number of electrons in that element. Then you have to simply arrange the electrons taking into account the maximum number that can be held in each shell. Remember, two electrons in the first shell, eight electrons in the second shell and eight electrons in the third shell. So electron configuration can be shown using shell diagrams or notation where we simply write the number of electrons in each shell and separate them by commas. For example, let's consider sodium or Na with an atomic number of 11. Its electron configuration in notation form would be two, eight, one.

[3:00]Atoms become ions because they want a stable full outer shell of electrons. This can be done in two ways, by giving away or taking in electrons. Let's take sodium as an example. Sodium has an atomic number of 11. So its electron configuration is two, eight, one. This means it has one electron in its outer shell. To achieve a full outer shell, sodium can do one of two things. It can lose one electron, so if it gets rid of that single electron in its outer shell, it will be left with eight, which is a full outer shell. Or it can gain seven electrons, so alternatively, it could gain seven more electrons to its outer shell. But that's harder and less likely. Now, consider which option is simpler. Is it losing one electron or gaining seven electrons? Losing one is much easier, so that's what happens. Sodium gives away its extra electron to achieve a full outer shell. We will learn this process in more detail later. So when a sodium atom which contains 11 electrons becomes a sodium ion, it loses one electron, leaving it with 10 electrons. So the electron configuration of a sodium ion is two, eight.

[4:52]Here's a periodic table. Notice that it's organized into columns and rows. The columns are numbered from one to eight using Roman numerals, and we call them groups. The rows are numbered from one to seven and we call them periods.

[5:17]Group eight, known as the noble gases, have full outer shells. Let's look at two examples from Group eight, helium and argon. As you can see, the number of electrons in these elements fills up their electron shells to the maximum, making them highly stable. Noble gases do not react with other elements because they are satisfied with their electron configuration.

[5:54]The number of outer shell electrons is equal to the group number in Groups I to VII. Here are some examples. Potassium has one electron in its outermost shell, so it is in Group one. Aluminum has three electrons in its outermost shell, so it is in Group three. And fluorine has seven electrons in its outermost shell, so it is in Group seven.

[6:32]The number of occupied electron shells is equal to the period number. Examples are, Potassium has four shells, so it's in Period four. Aluminum has three shells, so it is in Period three. And fluorine has two shells, so it is in Period two.

[7:02]Next, Isotopes. Isotopes are different atoms of the same element that have the same number of protons but different numbers of neutrons. So isotopes have the same atomic number since they have the same number of protons but different mass numbers due to the different number of neutrons. Let's look at an example of two carbon isotopes. C represents the element Carbon. This is Carbon 12 and this is Carbon 14. 12 and 14 are the mass numbers indicating the sum of protons and neutrons in the nucleus. Both isotopes have six protons.

[7:59]But Carbon 12 has six neutrons and Carbon 14 has eight neutrons.

[8:10]Ions are atoms with a net electric charge due to the gain or loss of electrons. Here is an example of a symbol for an ion.

[8:24]This is the chemical symbol of the element, Chlorine in this case. This is the mass number, for this example, it's the mass number of a Chlorine 35 isotope. And this is the charge of the ion. It may be represented as a superscript or as a signed number. For this example, the ion has a charge of minus one.

[8:56]Isotopes of the same element have the same chemical properties because they have the same number of electrons and therefore the same electronic configuration. So even though the mass numbers of isotopes are different, they will behave alike in chemical reactions because they share the same electronic configuration.

[9:26]The relative atomic mass, symbol A R of an element is a number you find on the periodic table right next to the element's name and the atomic number. The Relative Atomic Mass (Ar) is the average mass of all the isotopes of an element compared to 1/12th of the mass of a carbon-12 atom. The fixed mass of a carbon-12 atom is like a standard reference point for comparing the masses of all other atoms. We figure out relative atomic mass by looking at the relative masses of each isotope and how common each isotope is, that is, its abundance. This helps us understand an element's average atomic mass. So relative atomic mass can be calculated with this formula: percentage of isotope one multiplied by mass number of isotope one plus percentage of isotope two multiplied by mass number of isotope two divided by 100. This is the case when we are calculating the relative atomic mass of two isotopes. If we are calculating the relative atomic mass of say three isotopes, we will include that as well.

[11:03]Let's take a look at boron as an example, an element with two main isotopes. One isotope, boron 10, makes up about 20% of all boron atoms. The other isotope, boron 11, is more abundant, accounting for roughly 80% of boron atoms. Now, let's calculate the relative atomic mass of boron. For boron 10, abundance of 20% multiplied by its mass number of 10, plus for boron 11, abundance of 80% multiplied by its mass number of 11, divided by 100.

[11:54]So we'll get approximately 10.8, which is the relative atomic mass of boron. In the periodic table, the relative atomic mass of boron is mentioned as 11. That is because it has been rounded to a whole number.

[12:16]That concludes part two of Topic Two, Atoms, Elements and Compounds. Are you enjoying our videos? Are they helping you? Here is a way you can show your appreciation and support our continued efforts. You may use YouTube Super Thanks to send us thanks. Thank you. Hope this video helped you. Please share your thoughts and suggestions in the comment section. Thank you for watching and please don't forget to subscribe to IGCSE Study Buddy for more revision videos. Bye bye.