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This book is a collection of lies we taught to our Year 12 Chemistry students in their graduation year.
The lies include well-meaning simplifications of the truth, mistakes in the textbook, and, in a few extreme cases, blatant falsehoods.
This book isn’t a criticism of the VCE Chemistry course at all. In fact, I wrote this book to demonstrate the overwhelming complexity of Chemistry and the consequential need to make appropriate omissions and generalisations during our teaching as we tailor our lessons to the appropriate year level of students.
Rules taught as true usually work 90% of the time in this subject. Chemistry has rules, exceptions, exceptions to exceptions and so on. You’ll peel pack these layers of rules and exceptions like an onion until you reach the core, where you’ll find physics and specialist maths.
Click here to download We Lied to You (2019 edition).
Inspired by the formula booklets used by VCE Physics and VCE Maths Methods, here’s an 8-page Chemistry formula booklet you can use for your Year 11 and 12 Chemistry assignments. This custom-made booklet is a collection of reliable formulae that I have been using to answer VCE Chemistry questions while teaching and tutoring around Melbourne.
There are 76 formulae on 8 pages. At least 10 of these formulae aren’t in the three main chemistry textbooks. Orders are shipped in A4-sized booklet that resembles the VCAA Data Booklet.
Orders from schools, students and tutors are all welcome. Price includes free international delivery and a 10% voucher for the T-shirt store.
James Kennedy achieved outstanding A-level results in 2006 in Maths, Chemistry, Physics and Biology. Those excellent grades (which equate to an ATAR of 99+) earned him a BA (Hons) degree and a Masters degree in Natural Sciences from the University of Cambridge.
Shortcut formulae were just one of the techniques James used to pass his A-level exams and get into Cambridge. Along with structured revision, revision guides, practice papers and study notes on wall-cards, James used shortcut formulae to save precious time in the examination hall. You can get your own copy of these original shortcut formulae – revised and updated for the 2017-2021 VCE Chemistry course – for just $55 including free international shipping. Click here to get your copy.
Here’s an preview of the inside
VCE Chemistry Formula Booklet, $55. Free, Fast Delivery Included.
Oxygen from Theodore Gray’s amazing book, The Elements
This post concludes the Periodic Table Smoothie experiment.
Recall that we’ve just finished adding one mole of nitrogen gas and created a bizarre boron polymer at the bottom of our vessel. The temperature was 350 °C and the pressure in our vessel was 891 kPa.
Today, we’re going to add 1.00 mole of oxygen gas, stand back and observe.
Nothing happens.
This is disappointing news.
Many of the substances in our vessel react (more accurately, explode) in the presence of oxygen but the ignition temperature for all of those explosions to take place is at least 500 °C. The temperature of our vessel is set at just 350 °C. At this temperature, nothing would actually happen.
There’s not enough activation energy to break bonds in the reactant particles in order to get the reaction started. We call this activation energy (EA) in chemistry. If we were to add a source of excessive heat (e.g. a matchstick), the vessel would explode.
Should we heat up the vessel to 500 °C and blow up the experiment right here?
If we did, the following reactions would happen:
Enough of these reactions – particularly the first three – are sufficiently exothermic to trigger a chain reaction – at least up to the reaction of oxygen with beryllium carbide. The vessel would bang, explode, and shatter. The helium would float away, dangerous lithium amide would fly out sideways, and polyborazine powder, whatever that is, would land on the floor.
Let’s not ignite our experiment – not yet.
Conclusion after adding 1.00 mole of oxygen gas
Substance
Amount in mol
He(g)
1.000
Be(s)
0.514
LiH(s)
0.000
Li2C2(s)
0.272
B2H6(g)
0.000
Be2C(s)
0.175
H2(g)
0.007
BeC2(s)
0.136
CH4(g)
0.009
N2(g)
0.552
NH3(g)
0.154
LiNH2(s)
0.277
polyborazine
12.194 grams
Pressure: 891 kPa (higher than before due to the addition of nitrogen gas) Temperature: 350 °C (vessel is still being maintained at constant temperature)
Oxygen was relatively uneventful. Let’s add fluorine and see what happens.
Let’s add fluorine gas
Elements by Theodore Gray
The following three reactions would all occur as 1.00 mole of fluorine gas is added:
These two products are quite interesting:
HF, hydrogen fluoride, an aqueous solution of which was used by Breaking Bad’s Walter White to dissolve evidence (his victims)
NF3, nitrogen trifluoride, is used as an etching agent when making printed circuit boards (PCBs)
Let’s add neon gas
Elements by Theodore Gray
When 1.00 mole of neon gas is added, the total pressure inside the vessel increases but no reaction occurs. The concentrations of all the other gases present are unaffected.
The End
That concludes our Periodic Table Smoothie experiment. The most interesting conclusion was the discovery of polyborazine, the bizarre solid that collected at the bottom of the vessel.
Also of interest was how easily we created ammonia, one of the simplest of biological compounds, just by mixing elements together. Could the compounds necessary for life be so easy to create that their existence is an inevitable consequence of the Big Bang? Is life inevitable? If the Big Bang were to happen all over again, would life occur? And would it look any different?
This book contains 50 lies taught in the VCE Chemistry course.
These lies include well-meaning simplifications of the truth, mistakes in the textbook, and, in a few extreme cases, blatant falsehoods.
This book isn’t a criticism of the VCE Chemistry course at all. In fact, I just want to highlight the sheer complexity of Chemistry and the need to make sweeping generalisations at every level so it can be comprehensible to our students. This is a legitimate practice called constructivism in pedagogical circles. (Look that up.)
Many of these ‘lies’ taught at VCE level will be debunked by your first-year chemistry lecturers at university.
Here’s a preview of some of the lies mentioned in the book. Check out all 50 by clicking the download link at the bottom of the page.
The content you’re learning now is probably not as true as it seems. Chemistry is a set of models that explain the macro level sometimes at the expense of detail. The more you study Chemistry, the more precise these models become, and they’ll gradually enlighten you with a newfound clarity about the inner workings of our universe. It’s profound.
Rules taught as ‘true’ usually work 90% of the time in this subject. Chemistry has rules, exceptions, exceptions to exceptions, and exceptions to those – you’ll need to peel pack these layers of rules and exceptions like an onion until you reach the core, where you’ll find Physics and Specialist Maths.
Enjoy this book. I hope it emboldens you to question everything you’re told, and encourages you to read beyond the courses you’re taught in school.
Click to download REDOX RULES posters for VCE Chemistry
What’s redox? We never learned that!
Yes, you did. I use the term “redox” to refer to all of the following chapters in Heinemann Chemistry 2, which you will have learned at the end of Term 3 (September).
Chapter 26: Redox (revision of Year 11)
Chapter 27: Galvanic Cells
Chapter 28: Electrolytic Cells
Don’t underestimate redox
The VCAA has consistently used redox to discriminate which schools and students have the self-discipline required to keep studying at the end of the year. Studies show that redox is taught at a time when student motivation is at its minimum: energy levels are low, emotions are high, and graduation is just over the horizon. Many schools and students gloss over these topics because they’re running out of time, any many students think they’ve grasped the topic – when they’ve actually grasped misconceptions instead.
Notice how chapters 26, 27 and 28 are consistently the most difficult andthe most frequently asked? Click to download PDF version
Here are some popular redox lies (misconceptions)
LIE #1: The polarities switch during recharge Nope. The polarities never switch. It’s the labels of ‘anode’ and ‘cathode’ that switch because the electrons are flowing the other way through the external circuit. Polarity is permanent.
LIE #2: Hydrogen fuel cells don’t emit any greenhouse gases Wrong. They emit H2O, which is a powerful greenhouse gas. If you don’t believe that the VCAA can be this pedantic, think again. Read their 2015 Examiners Report here.
LIE #3: Each mole of electrons forms 1 mol Ag, 2 mol Cu or 3 mol Al in a cell Wrong again. If you look at the half-equations, you’ll see that each mole of electrons actually forms 1 mol Ag, 1⁄2 mol Cu or 1⁄3 mol Al. That’s why I teach “1, 1⁄2 and 1⁄3 moles” instead of the typical “1, 2, 3 moles” rule.
LIE #4: Temperature increases the rate of reaction in electroplating
Wrong! Remember that Faraday’s first law states that m ∝ Q. Because Q = I×t, only those two things – current and time – can affect the mass deposited at the cathode.
LIE #5: Electrons always leave the anode and go towards the cathode Wrong again. Electrons go RACO: to see what that means, download the posters above. This question appears in recent versions of Chemistry Checkpoints. Give it a try.
LIE #6: The cathode is always positive Ask your teacher.
LIE #7: Ions flow one way in the salt bridge
Nope. Anions always migrate to the anode; and cations always migrate to the cathode.
LIE #8: KOHES always works for balancing half-equations
KOHES only works for cells with acidic electrolytes. For cells with alkaline electrolytes, which sometimes appear in VCAA papers despite not being in the study design (see page 46 here), you’ll need to use KOHES(OH). Here’s KOHES(OH) explained:
Do KOHES as normal
Add the same number of OH–(aq) ions to each side of the half-equation to balance out the H+(aq)
Cancel and simplify. Remember that H+(aq) + OH–(aq) makes H2O(l). Remember also to cancel out any remaining H2O(l).
LIE #9: I can balance an unbalanced redox equation by putting numbers in the equation Don’t be fooled by this one! The ONLY way to balance an unbalanced redox equation successfully is to do the following:
Separate it into two half equations
Balance them using KOHES or KOHES(OH) as appropriate
Multiply them and recombine
Cancel and simplify
Done!
That’s a lot of work but it’s the only way to do it successfully. If you try to ‘cheat’ by just writing numbers (molar coefficients) in front of the reactants and products, you’ll find that the charges don’t add up, and you’ll get zero marks for the question.
LIE #10: I can break up polyatomic ions to make balancing half-equations easier
Nope! You’re only allowed to separate aqueous species in a half equation or an ionic equation. Because the Mn and O are actually bonded together in a polyatomic ion, you’ll need to write this:
If in doubt, keep it intact and it’ll cancel out by the end if it’s a spectator ion.
LIE #11: The two reactants that are closest together on the electrochemical series react Not always true. Use SOC SRA instead, which is explained in the posters above. Still struggling? Ask your teacher or tutor for help.
LIE #12: Oxidants are all on the top of the electrochemical series They’re actually on the left, and all the reductants can be found on the right side of each half equation in the electrochemical series. There is no top/bottom divide on the electrochemical series: only a left/right divide of oxidants/reductants.
Decorate your school/bedroom/hallway
Surround yourselves with truthful redox revision using these 17 free Redox posters. I’ve had these up around the whiteboard for a few weeks now – they’re a constant reminder to students that redox has many ideas that are always true.
One more tip: print and laminate an electrochemical series (available here) so you can annotate it during dozens of practice dozens without wasting paper. Good luck!
James Kennedy will explore the rise of chemophobia, an irrational fear of compounds perceived as ‘synthetic’, and the damage it can cause in this interactive webinar. We’ll examine its evolutionary roots, the factors keeping it alive today and how to fight chemophobia successfully.
What You Will Learn
Origins of chemophobia as an irrational psychological quirk
Chemistry teachers, Walter White, materialism and advertisements are all fuelling chemophobia today
Fighting chemophobia needs to be positive, respectful, multifaceted, and good for consumers
‘Boron’ page from Theodore Gray’s book, The Elements
Boron is a metalloid: an intermediate between the metals and non-metals. It exists in many polymorphs (different crystal lattice structures), some of which exhibit more metallic character than others. Metallic boron is non-toxic, extremely hard and has a very high melting point: only 11 elements have a higher melting point than boron.
British scientist Sir Humphrey Davy described boron thus:
“[Boron is] of the darkest shades of olive. It is opake[sic], very friable, and its powder does not scratch glass. If heated in the atmosphere, it takes fire at a temperature below the boiling point of olive oil, and burns with a red light and with scintillations like charcoal” – Sir Humphrey Davy in 1809
Initial condition
Before we add the 1.00 mol of boron into our reaction vessel, we need to recall what’s already in there from our experiments so far:
H2(g): 0.70 mol
He(g): 1.00 mol
Li(s): 0.40 mol
LiH(s): 0.60 mol
Be(s): 1.00 mol
The temperature of our vessel is 99 °C and the pressure of the gaseous phase is 525.5 kPa.
Now, let’s add our 1.00 mol of boron powder.
Which reactions take place?
Boron reacts with hydrogen gas to produce a colourless gas called borane, BH3(g), according to the following equation[1]:
Boron also reacts with lithium in very complex ways. If we heat the vessel up to 350 °C, we’d expect to see the formation of a boron-lithium system with chemical formula B3Li according to this equation[2]:
Notice that now we’ve heated up our vessel to 350 °C to allow this reaction to happen, the lithium at the bottom of the vessel has melted.
Boron reacts with lithium hydride as well, but only at temperatures around 688 °C. With our vessel’s temperature set at 350 °C, we won’t observe this particular reaction in our experiment.[3]
Some allotropes of boron – in particular, the alpha allotrope that was discovered in 1958 – is capable of reacting with beryllium to form BeB12. Because we’re using beryllium powder, which has semi-random symmetry, we won’t see any BeB12 forming in our vessel. Alpha-boron only exists at pressures higher than around 3500 kPa. At our moderate pressure of only 525.5 kPa, powdered (semi-random) boron will prevail and no BeB12 will form.[4]
For simplicity’s sake, let’s assume that the two reactions above take place with equal preference.
Boron powder reacts with hydrogen gas
Let’s do an ice table to find out how much borane we make.
A quick n/ratio calculation shows us that the hydrogen gas is limiting in this reaction:
We can expect all of the hydrogen gas to react with the boron powder.
units are mol
2 B
3 H2
2 BH3
I
0.50
0.70
0
C
-0.466
-0.70
+0.466
E
0.0333
0
0.466
Borane is very unstable as BH3, and it would probably dimerise into B2H6(g). This is still a gas at 350 °C and is much more stable than BH3. For the rest of this experiment we’ll assume that our 0.466 mol of BH3 has dimerised completely into 0.233 mol of B2H6.
Boron powder reacts with lithium
With the molar ratios present in our vessel, at 350 °C, we’d expect to witness the formation of a boron-lithium system, with chemical formula B3Li.
A quick n/ratio calculation shows that in this reaction, the boron powder is limiting.
All of the remaining boron therefore reacts with lithium. To calculate exactly how much B3Li we’ve created, let’s do another ice table:
units are mol
3 B
Li
B3Li
I
0.533
0.40
0
C
-0.533
-0.178
+0.178
E
0
0.222
0.178
What’s in our vessel after adding boron?
We have the following gas mixture in our vessel:
Helium gas, He(g): 1.00 mol
Helium is an inert noble gas that will probably remain in the vessel until the end of the experiment. It’s used in party balloons.
Borane gas, B2H6(g): 0.233 mol
We made this today. Borane is used in the synthesis of organic chemicals via a process called hydroboration. An example of hydroboration is shown below.
At the bottom of the vessel, there’s a sludge, which contains the following liquids and solids:
Molten lithium, Li(l): 0.22 mol
Lithium is used in the production of ceramics, batteries, grease, pharmaceuticals and many other applications. We’ve got 0.22 moles of lithium, which is about 1.5 grams.
Beryllium powder, Be(s): 1.00 mol
Beryllium is used as an alloying agent in producing beryllium copper, which is used in springs, electrical contacts, spot-welding electrodes, and non-sparking tools.
Lithium hydride, LiH(s): 0.60 mol
Lithium hydride is used in shielding nuclear reactors and also has the potential to store hydrogen gas in vehicles. Lithium hydride is highly reactive with water.
Boron-lithium system, B3Li(s): 0.178 mol
We made this today… but what is it? Not much is known about this compound – in fact, it doesn’t even have a name other than “boron-lithium system, B3Li”. It’ll probably decompose eventually in our experiment – maybe when we alter the pressure or temperature of the vessel at some later stage. We’ll need to keep an eye on this one.
The original H2(g) and B(s) have been reacted completely in our experiment.
What’s the pressure in our vessel now?
At the end of our reaction, the temperature of our vessel is still set at 350 °C and the pressure of the gaseous phase inside the vessel can be calculated to be a moderate 638 kPa as follows:
*It should also be noted that some evidence exists for a reaction between LiH and BH3, forming Li(BH4). The reaction seems to take place stepwise with increasing temperature. A quick read of this paper suggests that in our vessel, which is at 350 °C, any Li(BH4) formed would actually break back down into boron powder and hydrogen gas, which would in turn react with each other and with lithium metal to form BH3 and LiH again. The net result would be a negligible net gain of LiH and a negligible net loss of boron powder. We will continue calculating this Periodic Table Smoothie under the assumption that if any Li(BH4) forms, it breaks down before we add the next element, and the overall effect on our system is negligible.
**Li(BH4) is an interesting compound: it’s been touted as a potential means of storing hydrogen gas in vehicles – it’s safer and releases hydrogen more readily than LiH, which was mentioned above.[5]
Next week, we’ll add element number 6, carbon, and see what happens.
Yesterday, I was wondering what would happen if we mixed the entire periodic table of elements together in a blender. Unsurprisingly, it would explode, scattering radioactive dust and debris for miles around in a red-hot fireball formed from the simultaneous fission of the entire seventh row. The periodic table would only need to be the size of a matchbox in order for this explosion to happen.
Calculating exactly what would happen would be incredibly difficult. There are so many simultaneous reactions – including nuclear reactions – taking place that it’s almost impossible to predict the outcome in any more detail than “KABOOM”.
Making a real Periodic Table Smoothie would be prohibitively expensive. You’d need 118 particle accelerators (costing $1 billion each) all pointing at the same target just to get single atoms of each element to collide at the same time. This is even more difficult than it sounds: those elements near the bottom of the periodic table (numbers 105 and above) are so unstable that they’d break down before they even reach the target. There are massive financial and physical challenges to mixing an entire periodic table up in a blender.
Instead of adding all the elements at the same time, I’ll be adding one element each week to an imaginary 10-litre vessel and documenting – as a theoretical exercise – what happens. Ultimately, we all know it’s going to explode at some point. But when will it do that? How many elements are we able to add before it finally explodes? Will we create anything interesting along the way?
This very idea was floated on Reddit’s AskScience forum in 2013 but nobody actually figured out (seriously) what would happen.
How much of a chill will these ice cubes give to a bucket of hot water?Today, we’re going to answer the following question:
When 200 grams of ice is added to a bucket containing 1.00 litre of hot water, what’s the final temperature of the water?
To answer the question, we’re going to need to make some assumptions. We’ll take 1.000 litre of pure water at 80.00°C and add 200.0 g of ice (at -10.00°C) to it. What’s the final temperature of the water?
Part 1: Heat transfer method
The following equation can calculate the temperature at thermal equilibrium of any number of objects in thermal contact.
I love this equation because it’s several lines of maths shorter than the version taught in school. With this equation, you don’t even need to convert the temperatures into kelvin. Celsius works just fine.
Let’s set up the equation so that the addition series contains the variables in the question.
Now, let’s substitute the gives values into the equation. The specific heat capacity of water is 4200 J kg-1 K-1, and that of ice is 2100 J kg-1 K-1.
Great! Adding 200.0 g of ice to 1.000 L of water decreases the temperature from 80.00°C to 71.80°C.
But we’ve forgotten something. The ice will melt as soon as it hits the hot water. Since melting is an endothermic process, heat energy from the water will actually be absorbed, thus reducing the final temperature even further.
Part 2: Let’s take into account the fact that the ice melts!
Remember our formula from part 1.
The amount of energy required to melt ice can be calculated using the latent heat equation:
Removing that amount of heat energy from the system results in the following equation:
Great! Now, we’ve calculated that the final temperature of the water would be 57.36°C after the addition of the ice. That’s equal to 330.5 kelvin, which will be useful later.
However, we’ve forgotten to take something else into account: how much heat will be lost as radiation from the surface of the bucket?
Part 3: What’s the rate of heat loss from the bucket by radiation?
The rate of heat lost by radiation can be calculated by using the Stefan-Boltzmann equation, below.
P is the rate at which heat energy is radiated from the surface of the bucket in watts. Emissivity, e, of water is 0.95, and the surface area, A, should be around 0.0707 m2 for a one-litre bucket. Calculation of A is shown below. Assuming that the radius of the surface of the bucket is 6cm:
Plugging that value into the equation, we can find P. We’ll assume that the experiment is being conducted at room temperature and the temperature of the surroundings is 20.00°C (29.03 K).
This means that 2.928 joules of energy are emitted from the surface of the bucket every second. Ten minutes later, the bucket would have lost 1756.8 joules of energy due to radiation from the surface. But what about emission of radiation from the sides of the bucket?
Let’s say that our bucket is made from highly polished aluminium (which has emissivity 0.035) and it holds exactly 1.2 litres of water. We need to calculate the dimensions of the bucket.
Assuming it has straight sides (i.e. it’s a cylinder), the bucket had volume equal to the following formula:
The surface area of our bucket (excluding the open surface at the top) is:
The rate of energy radiation from the sides would therefore be:
It’s interesting to note how very little radiation is emitted from the shiny aluminium bucket, while lots more radiation is emitted from the surface of the water. This is because relatively ‘dark’ water has a much higher emissivity than shiny aluminium. Total emission from the bucket is therefore:
After ten minutes, the bucket would have lost the following amount of energy:
Let’s factor this amount of energy loss into our final temperature equation.
Not much energy is lost via radiation! Finally, let’s find the peak wavelength of the radiation emitted by the object using Wien’s law.
Part 4: What’s the wavelength of the radiation being emitted by the bucket?
Here’s Wien’s law from Unit 1 Physics…
The radiation emitted from the resulting bucket of water lies firmly in the infra-red part of the electromagnetic spectrum. The bucket would be clearly visible on an infra-red camera!
Next week, we’ll begin a new a Chemistry-themed project called Periodic Table Smoothie. More next week.
Chad Jones works at Intel Corp. in Utah, USA. He’s the founder and chief science writer for The Collapsed Wavefunction, a science advocacy podcast featuring episodes on science instruction, science in popular culture, and current science news items.
In 2016, Chad’s launched his latest venture in chemistry outreach with a fantastic new podcast called Chemical Dependence. In each of the podcast’s punchy, 5-minute episodes, Chad explores interesting chemical compounds and how they’re used in society. The podcast is a great source of interesting facts to liven up any chemistry lesson. All Chemistry teachers should subscribe!
He’s even teamed up with Andy Brunning from Compound Interest for his latest episode on pipeline. Check it out here.
Check out all the episodes and subscribe to the podcast on iTunes here. Support the podcast via Patreon here.
A new survey by Stop Procrastinating shows that social media is the biggest distraction students face while studying
The leading internet blocker, Stop Procrastinating, has announced that 64% of US students have cited online distractions such as social media as a hindrance to their productivity. Facebook, Twitter, Snapchat, shopping websites and YouTube were among the sites that students found the most distracting.
Fear of Missing Out (FOMO)
Nearly all of the students who responded in the survey referred to a ‘fear of missing out’ (FOMO), which is the anxiety that people experience when they believe that important events are happening without them. The anxiety arises from a perceived decrease in ‘popularity’ if they’re not up-to-date with the latest happenings in their social circle. Teenagers are particularly susceptible to FOMO, and 24-hour social media feeds such as Facebook and Twitter are exacerbating the problem. Students are constantly checking their social media feeds (sometimes a few hundred times per day) in order to keep up with the latest drivel happenings.
Interestingly, first year university students were the most affected. It’s possible that in first year (sometimes called “freshman year”), people’s social circles haven’t quite cemented since the upheaval of leaving high school. People are therefore more anxious and fear missing out on new friendships and events… so they gravitate towards social media.
Almost half students surveyed admitted to losing an hour each day to social media. Common Sense Media estimates the real figure (including traditional media such as TV) is more like 9 hours per day. That’s a lot of screen time, and it’s affecting students’ social lives, their grades and their sleep.
Over half of the respondents said they’d been stopped from writing an essay because they felt compelled to check social media at some point. Any issue that’s stopping half of our students from writing essays (or concentrating for any extended period of time) needs to be addressed urgently.
This problem needs to be addressed urgently
The level of distraction today is unprecedented. We all carry televisions and music players in our pockets. I got in touch with Tim Rollins, the director of Stop Procrastinating, who said:
“We have made Stop Procrastinating free today in order help students to beat their Internet distractions and boost their performance in their studies. The Internet, social media, emails are pervasive and eating into our quality time. We need urgently to put ourselves back in control.” – Tim Rollins
Software is one of the tools that can help students get the lasting willpower they need to overcome FOMO and get back into studying. Here are my tips for eliminating distractions while studying.
Tips for distraction-free studying
Delete all the Facebook apps from your phone
Study with your phone in aeroplane mode
When using your desktop, use the Stop Procrastinating app to limit your access to social media sites.
Study without music. All the research says it doesn’t help.
Don’t eat and study at the same time.
Drink only water while you’re studying.
Sit upright while studying: don’t study laying in bed or leaning back on the couch.
Have a goal for each study session. Write it down and work until you’ve completed it (e.g. make notes on all 6 types of acid/base chemical reactions with examples)
Study in a location that you never use for relaxation… the library is a great choice. Most students can’t study in their bedroom because they usually relax there.
Limit the number of Facebook friends to 30. Delete all the others: I understand this takes some courage, but you probably don’t know them anyway! Their unimportant updates distract you from studying.
Stop Procrastinating is an Internet blocking and productivity application compatible with OS X and Windows. It allows users the option to block the Internet for a period of time in three ways, depending on how much self-discipline they have.
“The Blue Marble” is a famous photograph of the Earth taken on December 7, 1972, by the crew of the Apollo 17 spacecraft en route to the Moon.
The rise of the environmental movement is most often attributed to the publication of Rachel Carson’s Silent Spring in 1962, which demonised chemicals as it introduced them to the public:
“Chemicals are the sinister and little-recognised partners of radiation entering into living organisms, passing from one to another in a chain of poisoning and death” – Rachel Carson’s Silent Spring, 1962
Later that decade, the Apollo missions and the six moon landings between 1969 and 1972 gave us a new perspective of planet Earth that was so profound that we felt a sudden compulsion to protect its natural beauty. Watch Neil deGrasse Tyson argue this point below.
In 1970, we are still going to the moon, we are still going until 1972, so watch these sequence of events. In 1970, the comprehensive Clean Air Act is passed… Earth Day was birthed in March 1970. The EPA was founded in 1970… Doctors Without Borders was founded in 1971… DDT gets banned in 1972, and we are still going to the moon. We’re still looking back at Earth. The clean water act 1971, 1972 the endangered species act, the catalytic converted gets put in in 1973, and unleaded gas gets introduced in 1973… That is space operating on our culture and you cannot even put a price on that. – Neil deGrasse Tyson in April 2012
Together, Rachel Carson and the Apollo missions made the public in Western countries quickly aware of the Earth and its natural beauty. Humans were portrayed as selfish destructors of a planet that was supposedly most ‘beautiful’ when in its ‘natural’ state. The field of toxicology was spawned in wake of this concern, and had the goal of analysing the toxicity of different chemicals on humans and the environment. As the first edition of Human and Experimental Toxicology stated:
“Politicians cannot be expected to come to rational and acceptable decisions without adequate impartial and objective information, and toxicologists have grave responsibilities to produce such information”. – Human and Experimental Toxicology
While the field of toxicology accumulated a wealth of scientific evidence about ‘chemicals’, this evidence largely hasn’t trickled down to the public and certainly hasn’t allayed their fears. There remains a lingering skepticism about chemicals, especially artificial chemicals, which some people still feel are more harmful than those found in nature.
Take the Think Dirty iOS app, for example, which gives cosmetic ingredients a safety rating out of 9. According to the app’s creators, “Fragrance” gets the worst possible rating (9), while “Natural Fragrance” gets the best rating (1). Black-and-white ‘natural’ vs ‘artificial’ decision-making such as this is completely unfounded and ignores toxicological evidence. This kind of thinking is misleading, has no scientific basis and sometimes causes consumers to make harmful conclusions – no matter how benign their intentions. (More on this in future posts.)
This simplistic thinking is a remnant of the environmental movement back in the 1970s: that ‘selfish’ humans were destroying a ‘pristine’ planet Earth. While the ‘natural/good’ vs ‘artificial/bad’ dichotomy was an effective solution to short-term environmental problems of the time, this black-and-white thinking is actually leading people to make bad decisions today. We can no longer assume that “natural” is always “best”: the issue is actually far more complex than that. Toxicological evidence needs to be made public and easy to digest so that consumers can make more enlightened decisions.
This post is part 1 of a weekly series on Chemophobia. More next week.
It’s been two years since I posted the All-Natural Banana. Motivation behind this poster was to dispel the myth that “natural = good” and “artificial = bad”. It’s been a very successful project. It’s spawned 11 more “Ingredients” posters, a successful clothing line, and has sold thousands of print copies worldwide via this website.
Online news portal io9then published a news story about the All-Natural Banana, which was followed in quick succession by articles in Vox, Forbes, Business Insider, the New York Timesand more.
From today onwards, you can download the original PDF artworks for free. They come with a Attribution-NonCommercial 4.0 International Creative Commons License, which means that you can share them, print them and modify them as much as you like for non-commercial purposes only.
I’ll be following this up with an article on the ‘Origins of Chemophobia’ next week. Subscribe to this website below or subscribe via my Apple News channel here.
James Kennedy 